الفهرس 
Electrochemical behavior of Mg and Mg alloys in differentOnly 14 pages are availabe for public view
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References
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Studies on Heavy Metals Removal from Polluted Water Using Some Organic and Inorganic Composite Materials
A Thesis Submitted by
Mohamed Ahmed Abdo Abd Allah Elamir
B.Sc in Chemistry
In Partial Fulfillment of the Requirements For
The Degree of master of Science (M.Sc.)
in
Chemistry (Inorganic and Analytical Chemistry)
to
Chemistry Departement
Faculty of Sience
Suez Canal University
Ismailia
(2019)
1. Introduction and Literature review
Heavy metals are metallic elements which have a comparatively high density compared to water. With the assumption that overweight and toxicity are interrelated, heavy metals also include metalloids, such as arsenic, that are able to motivate toxicity at low exposure level. Nowadays, the environmental pollution by these metals causes extreme ecological and global public health worry. Also, human exposure has risen intensely owing to an exponential increase of their use in various industrial, domestic, agricultural, and technological applications. Generally heavy metals sources involve industrial, geogenic, agricultural, pharmaceutical, domestic waste, and atmospheric sources. Environmental pollution is very notable in point source areas like mining, foundries and metal extraction, and other metal-based industrial operations [1-3].
Even though heavy metals are naturally occurring elements that are present throughout the earth’s crust, most environmental pollution and human exposure result from anthropogenic activities such as mining and smelting operations, industrial production and use, and domestic and agricultural use of metals and metal-containing compounds [3-5]. Environmental contamination can also occur through corrosion of metal, atmospheric deposition, soil erosion of metal ions and leaching of heavy metals, sediment resuspension, and metal evaporation from water resources to soil and groundwater. Natural phenomena such as weathering and volcanic eruptions have also been reported to significantly contribute to heavy metal pollution [1-3]. Industrial sources of heavy metals include metal processing in refineries, coal burning in power plants, petroleum combustion process, stations of nuclear power and high-tension lines, plastics, textiles, microelectronics, wood preservation, and paper-processing plants [6, 7]. It has been stated that metals such as iron, magnesium, manganese, chromium, molybdenum, cobalt, nickel, copper, selenium and zinc are necessary nutrients which are needed for various physiological and biochemical functions [1]. Insufficient supply of these micronutrients leads to a variety of deficiency diseases or syndromes.
Also, heavy metals are considered as trace elements owing to their presence in very small concentrations (ppb range to less than 10 ppm) in various environmental regions. Their bioavailability is affected by physical factors such as temperature, phase combination, adsorption, and sequestration. It is also influenced by chemical factors which affect speciation at thermodynamic equilibrium, kinetics of complexation, lipid solubility, and octanol/water partition coefficients. Biological factors, such as characteristics of species, trophic interactions, and biochemical/physiological adaptation, also play an important role [1, 8].
The essential heavy metals perform biochemical and physiological functions in plants and animals. They are important constituents of several key enzymes and play important roles in various redox reactions [1]. Copper, for example, serves as an essential cofactor for sundry oxidative stress-related enzymes including catalase, superoxide dismutase, peroxidase, cytochrome c oxidases, dopamine β-monooxygenase, monoamine oxidase, and ferroxidases [9-11]. Therefore, it is an essential nutrient that is incorporated into a number of metalloenzymes involved in hemoglobin formation, carbohydrate metabolism, catecholamine biosynthesis, and collagen cross-linking, elastin, and hair keratin. The ability of copper to cycle between an oxidized state, Cu(II), and reduced state, Cu(I), is used by cuproenzymes involved in reduction-oxidation reactions [9-11]. However, it is this characteristic of copper that also makes it potentially toxic as a result of the transitions between the two oxidation states Cu(II) and Cu(I) can lead to superoxide and hydroxyl radicals generation [9-12].
Additionally, too much exposure to copper causing cellular damage leading to Wilson disease in humans [11, 12]. As in the case of copper, several other essential elements are required for biologic functioning; however, an excess amount of this metals causes cellular and tissue damage leading to a variety of harmful effects and human diseases. For some elements including copper and chromium, a very small range of concentrations between useful and toxic effects is present . Other metals such as cadmium, antinomy, arsenic, aluminum, barium, beryllium, bismuth, gallium, germanium, gold, indium, lead, lithium, nickel, mercury, platinum, silver, strontium, tellurium, thallium, tin , titanium, vanadium, and uranium have no definite biological functions and are regarded as nonessential metals [1].
In biological systems, heavy metals have been reported to affect cellular organelles and components like lysosome, cell membrane, endoplasmic reticulum, nuclei, mitochondrial, and some enzymes involved in detoxification, damage repair, and metabolism [13]. Metal ions have been found to interact with cell components such as nuclear proteins and DNA, causing DNA damage and conformational changes that may lead to cell-cycle modulation, carcinogenesis, or apoptosis [13, 14]. Table (1) shows the main sources, health effects and permissible limits of various toxic heavy Metals.
Table (1): Sources, health effects and permissible limits of various toxic heavy Metals according to World Health Organization(WHO) [15, 16].
Metal
Source
Potential health effect potable water limits (ppm), WHO
Copper
Zinc
Mercury
Nickel
Cadmium
Arsenic
lead
Metal finishing industry
Electroplating
Metalliferous mining
Metal finishing industry
Electroplating, Fertilizers
Metalliferous mining
Agricultural material
Manures sewage sludge
Electronics
Waste disposal
landfill leachate
Metalliferous mining
Metal finishing industry
Electrodeposition
Manures sewage sludge
Alloys and steels
Metallurgical industries
Metalliferous mining
Agricultural materials
Fertilizers, waste disposal
Landfill leachate, electronics
Electronics
Metallurgical industries
Manures sewage
Specialist alloys and steels.
waste disposal
landfill leachate
Electronics, metallurgical
Industries
Nervous system irritation followed by depression
Liver damage, Wilson disease, insomnia
Phytotoxic, depression
Anemia, lethargy
Lack of muscular coordination
Abdominal pain
Increased thirst
Poisonous
Rheumatoid arthritis
Disturbs the cholesterol
diseases of the kidneys, circulatory system and nervous system
High conc. can cause DNA damage
Eczema of hands
Human carcinogen
High phytotoxicity
Damaging fauna
Kidney damage, renal disorder
Human carcinogen
Bronchitis, emphysema
Anemia
Acute effects in children
Skin manifestations,
visceral cancers, vascular disease
Immunotoxic
Modulation of co-receptor expression
Damage the fetal brain
diseases of the kidneys
circulatory system and
nervous system
2
3
0.001
0.02
0.03
0.01
0.01
1.1. `Mechanism of heavy metals in human being.
In past decades, several studies have been carried out to investigate the mechanism of toxicity with heavy metals [15] . Toxicity and carcinogenicity of heavy metals involve many mechanistic aspects, some of which are not clearly elucidated or understood [1].
Many studies have reported that the production of reactive oxygen species (ROS) and oxidative stress play an important and a key role in the toxicity and carcinogenicity of metals for example arsenic [17-19], cadmium [20] , chromium [21, 22] , lead [23, 24], and mercury [25, 26] . For the reason of their high degree of toxicity, these previous metals are among the priority metals that are of great public health significance. They are all systemic toxicants that are known to prompt multiple organ damage, even at lower levels of exposure. As per the US Environmental Protection Agency (US EPA) and the International Agency for Research on Cancer (IARC), these elements are also categorized as either “known” or “probable” human carcinogens based on epidemiologicalds and experimental studies showing a link between exposure and cancer incidence in humans and animals.
Oxidative stress is one of the major mechanisms behind metal toxicity [27]. The formation of large amounts of reactive oxygen species, such as superoxide anion (O2.-), hydrogen peroxide (H2O2), hydroxyl radical (HO.) and singlet oxygen (1O2), has been reported to promote the induction of oxidative stress. [15, 28] In fact, various studies connect heavy metals with oxidative DNA damage since these metals may reduce the level of the main antioxidant compounds in several animal tissues by inactivating enzymes and other antioxidant molecules [29]. In humans, oxidative stress is also responsible for various diseases, including cancer, Parkinson’s disease, Alzheimer’s disease, atherosclerosis, heart failure and myocardial infarction [30, 31].
Although Heavy metal-induced toxicity and carcinogenicity involve many mechanistic aspects, some of which are not clearly elucidated or understood. Each metal is known to have unique features and physicochemical properties that confer to its specific toxicological mechanisms of action [1].
1.1.1 BIOCHEMISTRY OF TOXICITY
The heavy metals poisoning effects are because of their interference with the normal body biochemistry in the normal metabolic processes. When this metals ingested, in the acid medium of the stomach, they are converted to their stable oxidation states (Zn2+, Pb2+, Cd2+, As2+, As3+, Hg2+ and Ag+) and combine with the body’s biomolecules such as proteins and enzymes to form stable and strong chemical bonds. The equations shown in Figure (1) display their reactions during bond formation with the sulphydryl groups (-SH) of cysteine and sulphur atoms of methionine (-SCH3) [32, 33].
Figure (1): interaction between metal with proteins and enzymes.
(A) = Intramolecular bonding; (B) = Intermolecular bonding; P = Protein; E = Enzyme; M = Metal
The metal groups or the hydrogen atoms in the above case are substituted by the poisoning metal and the enzyme is thus inhibited from functioning, whereas the protein–metal compound works as a substrate and can reacts with a metabolic enzyme. In the following scheme, equation C indicates the reaction of enzymes (E) with substrates (S) in either the lock-and-key pattern or the induced-fit pattern. In both cases, a substrate fits into an enzyme in a highly specific fashion, as aresult of enzyme chirality’s, to form an enzyme–substrate complex (E-S*) as follows [33] .
(E = Enzyme; S = Substrate; P = Product; * = Activated Complex)
While at the E-S, E–S* and E-P states, an enzyme cannot accommodate any other substrate till it is freed. Occasionally, the enzymes for an entire sequence coexist together in one multi-enzyme complex consisting of three or four enzymes. The product from one enzyme reacts with a second enzyme in a chain process, with the last enzyme yielding the final product as follows:
The final product (F) goes back to react with the first enzyme thereby inhibiting further reaction since it is not the starting material for the process. Hence, the enzyme E1 becomes incapable of accommodating any other substrate until F leaves and F can only leave if the body utilizes it. If the body cannot utilize the product formed from the heavy metal – protein substrate, there will be a permanent blockage of the enzyme E1, which then cannot initiate any other bio-reaction of its function. Therefore, the metal remains embedded in the tissue, and will result in bio-dysfunctions of various gravities. Furthermore, a metal ion in the body’s metallo-enzyme can be conveniently replaced by another metal ion of similar size. Thus Cd2+ can replace Zn2+ in some dehydrogenating enzymes, leading to cadmium toxicity. In the process of inhibition, the structure of a protein molecule can be mutilated to a bio-inactive form, and in the case of an enzyme can be completely destroyed. For example, toxic As3+ occurs in herbicide, fungicides and insecticides, and can attack –SH groups in enzymes to inhibit their bioactivities as shown below in Figure (2) [32, 33].
Figure (2): interactions between arsenic and enzyme.
The most toxic forms of these metals in their ionic species are the most stable oxidation states. For example, Cd2+, Pb2+, Hg2+, Ag+ and As3+. In their most stable oxidation states, they form very stable biotoxic compounds with the body’s bio-molecules, which become difficult to be dissociated, due to their bio-stabilities, during extraction from the body by medical detoxification therapy.
1.2. Heavy metals treatment techniques:
Heavy metal uptake from inorganic effluent and industrial wastewaters can be carried out by traditional treatment processes, such as, complexation, adsorption, coagulation, ion exchange, solvent extraction, chemical precipitation, electroplating, cementation, flotation and membrane separation. All of these processes may be physical, chemical or biological as shown in Figure (3), Some of these are illustrated in Figure (4) [34] . Different methods, such as chemical precipitations, conventional adsorption [35-37] , ion exchange [38], membrane separation techniques [39] and electro-remediation techniques are used usually for industrial wastewater treatment. Precipitation is most economical and hence widely used, but many industries still use chemical procedures for treatment of effluents due to economic considerations. The efficiency of the precipitation process is extremely decrease owing to the presence of complexing agents in wastewater , and this lead to incomplete processing and production of toxic sludge. Thus several new approaches have been studied to develop low cost effective and more efficient heavy metal adsorption techniques [40].
Biosorption is considered as a user-friendly with specific affinity, low cost and simple design so it is an effective separation and purification method for heavy metals disposal from industrial wastewater and it has been widely used for this purpose [41, 42].
Figure (3): Conventional technologies for heavy metal removal.
Sorption with sorbents made of agricultural or industrial by-products are used widely for heavy metals uptake from aqueous mediums because of their abundant availability, promising physical, low cost, and, surface and chemical characteristics [43]. Those materials and methods were widely discussed meeting their advantages.
Figure (4): Some conventional methods for metal removal.
1.2.1. Physico-chemical methods
Following methods have been used by various researchers for heavy metals uptake. Physical separation process are primarily applicable to particulate forms of metals, metal-bearing particle or discrete particles . Physical separation consists of flotation, mechanical screening, gravity concentration, hydrodynamic classification, magnetic separation, attrition scrubbing, and electrostatic separation, physical separation efficiency depends on several soil characteristics such as particulate shape, particle size distribution, moisture content, humic content, clay content, density between soil matrix ,metal contaminants and heterogeneity of soil matrix, and hydrophobic properties, magnetic properties of particle surface [40, 44].
The conventional chemical processes for heavy metals disposal from waste water contain many processes such as flotation, ion exchange, adsorption, electrochemical deposition and chemical precipitation. Factors which may limit the effectiveness and applicability of the chemical process are high content of clay/silt, calcite, humic, Ca and Fe, anions, heavy metals, or high buffering capacity [45].
A. Chemical Precipitation:
Chemical precipitation is one of the most widely used for removal of heavy metal from inorganic effluent in industry because of its simple operation[46]. These conventional chemical precipitation processes yield insoluble precipitates of heavy metals as hydroxide, carbonate, phosphate and sulfide. The mechanism of this process is depend on to produce insoluble metal precipitation by reacting dissolved metals in the solution and precipitant. In the precipitation method very fine particles are generated and chemical precipitants, flocculants and coagulation processes are used to increase their particle size in order to remove them as sludge [45, 46]. As soon as the metals precipitate and form solids, they can easily be removed, and metal with low concentrations, can be discharged. Removal percentage of metal ions in the solution may be improved to optimum by changing major parameters such as pH, initial concentration, ions charge, temperature,.…etc. Hydroxide treatment is the most commonly used precipitation technique owing to its relative simplicity, low cost of precipitant (lime), and ease of automatic pH control. The solubilities of the various metal hydroxides are minimized for pH in the range (8-11).
B. Coagulation and Flocculation:
The coagulation-flocculation mechanism is based on zeta potential (ζ) measurement as the standard to define the electrostatic interaction between coagulant-flocculant agents and pollutants [47] .Coagulation process is reduced the net surface charge of the colloidal particles to stabilize by electrostatic repulsion process [40]. Flocculation process continually increases the particle size to discrete particles through additional collisions and interaction with inorganic polymers formed by the organic polymers added [48]. The minute discrete particles are flocculated into larger particles, they can be removed or separated by filtration, floatation or straining. Sludge production, transfer of toxic compounds into solid phase and application of chemicals are main disadvantages of this process.
C. Electrochemical Treatments:
Electrolysis: Electrolytic recovery is one technology used for removing metals from waste water streams. This process uses electricity to pass a current through an aqueous metal-bearing solution containing a cathode plate and an insoluble anode. Electricity can be generated by movements of electrons from one element to another. Electrochemical process to treat wastewater containing heavy metals is to precipitate the heavy metals in a weak acidic or neutralized catholyte as hydroxides. Electrochemical treatments of wastewater involve electro-deposition, electro-coagulation, electro-flotation and electro-oxidation [49].
Electro-destabilization of colloids is called coagulation and precipitation by hydroxide formation to acceptable levels. It is the most common heavy metal precipitation process forming coagulants by electrolytic oxidation and destabilizing pollutants to form folc [50]. The electro-coagulation process the coagulant is generated in situ by electrolytic oxidation of an appropriate anode material. In this process, charged ionic metal species are removed from wastewater by allowing it reacting with anion present in the effluent. This process is characterized by reduced production of sludge, ease of operation and no requirement for chemical use.
However, chemical precipitation requires a large amount of chemicals to reduce metals to permissible limit for discharge. Other drawbacks are huge sludge production, poor settling, slow metal precipitation, long-term environmental impacts of sludge disposal , and the aggregation of metal precipitates [51]. It converts the aqueous pollution problem to a solid waste disposal problem without recovering the metal.
D. Ion Exchange:
Ion exchange can attract soluble ions from the liquid phase to the solid phase. It considered is the most widely used technique in water treatment industry. As a cost-effective method, ion exchange process usually involves convenient operations and materials with low-cost, and it has been demonstrated to be very effective for elemination heavy metals from aqueous mediums, specific for treating water with low heavy metals concentration [52, 53]. In this technique cations or anions containing special ion exchanger is used to eliminate metal ions from the solution. Commonly used ion exchangers are synthetic organic ion exchange resins. It can be used only low concentrated metal solution and this method is extremely sensitive with the pH of the aqueous phase.
Ion exchange resins are water-insoluble solid substances which can absorb negatively or positively charged ions from an electrolyte solution and release other ions with the same charges into the solution in an equivalent amount. The positively charged ions in cationic resins such as sodium and hydrogen ions are replaced with positively charged ions, such as, copper, zinc and lead ions, in the solutions. In a similar way, the negative ions in the resins such as hydroxyl and chloride ions can be exchanged by the negatively charged ions such as nitrate, chromate, sulfate, cyanide, and dissolved organic carbon (DOC).
E. Membrane Filtration:
Membrane filtration has received a great attention for the remidation of inorganic effluent. It is able to remove organic compounds, suspended solid, and inorganic pollutants such as heavy metals. Depending on the particle size that can be retained, different types of membrane filtration such as nanofiltration, ultrafiltration, and reverse osmosis can be utilized for heavy metal uptake from wastewater.
Ultrafiltration (UF) utilizes permeable membrane to separate heavy metals, macromolecules and suspended solids from inorganic solution on the basis of the pore size ranging from 5 to 20 nm and molecular weight of the isolating compounds (1000– 100,000 Da) [54]. Based on the membrane characteristics, UF can achieve more than 90% of removal efficiency with a metal concentration (10 - 112 mg/L) at pH ranging between 5 and 9.5 and at pressure (2–5 bar) . UF offers some advantages such as smaller space requirement and a lower driving force owing to its high packing density.
Polymer-supported ultrafiltration (PSU) process adds water soluble polymeric ligands to bind metal ions and form macromolecular complexes by generating a free targeted metal ions effluent [55]. PSU technology has the dvantages of the low-energy requirements involved in ultrafiltration, the reaction kinetics is very fast and higher separation selectivity of selective bonding agents in aqueous solution.
Another similar technique, complexation–ultrafiltration, confirms to be a promising alternative to technologies depend on ion exchange and precipitation. Using water-soluble metal-binding polymers in combination with ultrafiltration (UF) is a hybrid approach in order to concentrate selectively and to recover heavy metals in the solution. In the complexation – UF process cationic forms of heavy metals are first complexed by a macro-ligand in order that increasing their molecular weight with a size larger than the selected membrane pores [56, 57]. The advantages of complexation–filtration process iclude high selectivity of separation because of the use of a selective binding and low-energy requirements involved in these processes. Water-soluble polymeric ligands have shown to be powerful substances in order to separate trace metals from aqueous solutions and industrial wastewater through membrane processes.
Reverse osmosis (RO) is a separation process using pressure in which solution is forced through a membrane that keeps the solute on one side and allows the passage of the pure solvent to the other side. The membrane here is semi-permeable, meaning it allows the passage of solvent but not for metals. The reverse osmosis membranes have a dense barrier layer in the polymer matrix where most separation takes place. Reverse osmosis can remove many types of ions and molecules from solutions, including bacteria, and is used in both industrial processes. A diffusive mechanism is involved in reverse osmosis process, so that separation efficiency is dependent on pressure, concentration of the solute, , and water flux rate [58].
F. Electrodialysis:
Electrodialysis (ED) is a membrane separation uses electric potential to passe ionized species in solution through an ion exchange membrane. The membranes are plastic materials thin sheets with either cationic or anionic characteristics. When a solution containing ionic species passes through the compartments of the cell, the anions migrate toward the anode while the cations migrate toward the cathode, crossing the anion exchange and cation-exchange membranes [59]. A disadvantage of this process include membranes replacement and the corrosion process [60]. Using membranes with higher capacity of ion exchange resulted in better cell performance. Effects of temperature, flow rate,and voltage at different concentrations by using two types of commercial membranes, using a laboratory ED cell, on the removal of lead were studied [61]. The princible of Electrodialysis process is illustrated in Figure (5). Results show that increasing temperature and voltage improved cell performance and separation percentage decreased as the flow rate increasing. This provides advantages for the treatment of highly concentrated wastewater laden with heavy metals to recovery undesirable impurities from water.
Figure (5): Electrodialysis principles [62] CM – cation exchange membrane, D-dialute chamber, e1 and e2-electrode chambers, AM-anion exchange membrane and K-concentrate chamber
G. Adsorption:
Biosorption is another technique that can used for elimination of heavy metals from wastewater. Sorption process is defined as transfer of ions from solution phase to the solid phase, really describes a group of processes, which includes adsorption and precipitation reactions. Adsorption has become one of the alternative remidation techniques for wastewater. Basically, adsorption is a mass transfer process and substances bound by chemical and or physical interactions to solid surface [63-65]. All adsorption mechanisms are dependent on solid-liquid equilibrium and on mass transfer rates. A dsorption could be divided into the following types, depending on the types of intermolecular attractive forces [66, 67].
Physical adsorption:
It is a process in which binding of adsorbate on the surface of adsorbent as a result of Van der Waals forces of attraction or hydrogen bonding. Physical adsorption can only be occurred in the low temperature environment and under appropriate pH conditions.
Chemical adsorption:
A strong interaction arise from chemical reaction between the adsorbate and the adsorbent molecules is involved. This interaction produces new types of electronic bonds (Covalent and Ionic).
Mechanism of adsorption:
generally, the main steps involved in adsorption of contaminates on solid adsorbent are:
1.Transfer of the metal ion from the bulk of solution to the outer surface of the adsorbent.
2. Internal mass transfer by pore diffusion from outer surface of adsorbent to the inner surface of porous structure.
3. Adsorption of adsorbate onto the active sites of the adsorbent pores
4. The overall adsorption rate is determined by either intra particle diffusion or film formation or both as the last step of adsorption are very fast as compared to the other two steps.
The parameters which have been established for optimizing the use of adsorbent in wastewater treatment include [34]:
1. Nature of adsorbent and adsorbate.
2. Metal concentration.
3. pH and temperature of the aqueous solution.
4. Kinetics of adsorption.
5. Adsorption isotherm.
6. The time of contact.
Various low-cost adsorbents, derived from natural material, agricultural waste, industrial by-product, or modified biopolymers are found to be more promising and encouraging in heavy metal removal owing to various considerations as follow [36, 63].
(I)They are economical, (II) its metal selectivity, (III) they are regenerative, (IV) toxic production of sludge not present (V) metal recovery and (VI) its high effectiveness.
Using activated carbon in water and wastewater remidation has been directed towards organics removal [40] . Research efforts on removal of inorganics by activated carbon, specifically metallic ions, have been markedly limited [40] . selective adsorption by red mud [68], coal [69], photocatalyst beads [70], nano-particles[71], fertilizer industrial waste [72], biomass [73], activated sludge biomass [74], algae [75, 76] etc. has generated increasing excitement.
Industrial by-products such as fly ash [77] iron slags, waste iron [78] , hydrous titanium oxide [79, 80] ,can be chemically modified to enhance its removal performance for metal elimination from wastewater.
It was reported that [81, 82] for the disposal of heavy metals from industrial waste effluent has been focused on the use of agricultural by-products as adsorbents through biosorption process. New resources such as rice husk, coconut shell, pecan shells, rice straw, maize cob or husk, jackfruit, hazelnut shell, rice husk,…etc can be used as an adsorbent after chemical modification or conversion by heating into activated carbon or biochar for heavy metal uptake. They found that the maximum metal removal occurred by those biomass due to containing of cellulose, lignin, carbohydrate and silica in their adsorbent [83] .
Biopolymers are posse a number of different functional groups, such as amines and hydroxyls, which increase the efficiency of metal ion uptake [84] . They are widely use in industrially as they are able to lower the concentrations of transition metal ion to sub-part per billion concentrations. New polysaccharide-based-materials are described as biopolymer adsorbents (derived from chitosan, starch and chitin) for the elimination of heavy metals from the wastewater. The sorption mechanisms of polysaccharide-based-materials are complicated and depend on pH [84]. Also hydrogels, which are cross linked hydrophilic polymers, are widely used to purify wastewater. The removal is mainly governed by the water diffusion into the hydrogel, carrying the heavy metals inside especially in the nonexistence of strongly binding sites. Maximum binding capacity increases with higher pH because of polymerization/cross linking reaction.
1.2.2. Biological Methods:
Biological removal of heavy metals in wastewater involves the use of biological methods for the elimination of pollutants from wastewater. In this processes microorganisms play an important role of settling solids in the solution. Activated sludge, stabilization ponds, trickling filters are widely used for wastewater purification. Activated sludge is considered the most common option uses microorganisms in the treatment process to break down organic matter with agitation and aeration, and then allows solids to settle out. Bacteria-containing “activated sludge” is frequently re-circulated back to the aeration basin to increase organic decomposition rate. In biological systems, most of the research on heavy metals removal has been oriented towards the suspended growth activated sludge process. Trickling filters which consist beds of coarse media (often plastic or stones) 3-10 ft. deep help to grow microorganisms. Wastewater is sprayed into the air (aeration), then allowed to trickle through the media and microorganisms break down organic matters in the wastewater. The drain of trickling filters at the bottom and the wastewater is collected and then undergoes sedimentation. Lagoons or stabilization ponds are cheap, slow and relatively inefficient, biological method that can be used for different types of wastewater. They depend on the interaction of sunlight, microorganisms, algae, and oxygen [40].
1.3. Evaluation of heavy metals removal processes:
Although all the techniques of heavy metal wastewater treatment can be employed to eliminate heavy metals, they have their latent advantages and limitations. Table (2) indicates heavy metals uptake from aqueous solutions has been traditionally carried out by chemical precipitation because it is a simple process and cheap capital cost. However, chemical precipitation is ordinarily adapted for treating wastewater containing heavy metal ions with high concentration and it is ineffective with low metal ion concentration. Chemical precipitation is considered as not economical and can produce large sludge amount to be treated with great drawbacks [85].
Ion exchange has been commonly applied for the removal of heavy metal from wastewater. However, the resins of ion-exchange must be regenerated by chemical reagents when they are exhausted and the regeneration can cause serious secondary contamination. And it is expensive, particularly when treating a large amount of wastewater containing low concentration heavy metal, so that they cannot be used at large scale.
Adsorption is a common method for the uptake of heavy metals from low concentration aqueous solutions containing heavy metal. The activated carbon high cost limits its use in adsorption. Many varieties of adsorbents with low-cost have been developed and tested for heavy metal ions uptake. However, the efficiency of adsorption depends on the adsorbents type. Biosorptio of heavy metals from aqueous mediums is considered as new method that has proven very promising for the removal of heavy metal ions from wastewater [85].
The technology of membrane filtration can separate heavy metal ions with high efficiency, but its drawbacks such as high cost, low permeate flux process complexity, and membrane fouling have limited their use in heavy metal ions uptake.
Coagulation-flocculation technique can be employed for heavy metal wastewater remidation, the advantages of this method are dewatering and good sludge settling of the produced sludge. large sludge volume generation and chemical consumption are the limitations of this technique.
Table (2): The main advantages and disadvantages of the various physico-chemical methods for treatment of heavy metal in wastewater.
Treatment method
Target of removal
Advantages
Disadvantages
References
Chemical precipitation
Coagulation–flocculation
Dissolved air flotation
Ion exchange
Ultrafiltration
Nanofiltration
Reverse osmosis
Adsorption with new adsorbents
Heavy metals, divalent metals
Heavy metals and suspended solids.
Heavy metals and suspended solids
Dissolved compounds,
cations/anions
High molecular weight compounds (1000–10000 Da)
Hardness ions such as Ca(II) and Mg(II) and sulphate salts
Organic and inorganic compounds
Heavy metals
Low capital cost, simple operation
Shorter time to settle out suspended solids, improved sludge settling.
Low cost, shorter hydraulic retention time
No sludge production, less time consuming
Smaller space requirement
Lower pressure than RO (7–30 bar)
High rejection rate, able to withstand high temperature
Low-cost, easy operating conditions, having wide pH range, high metalbinding capacities
Sludge generation, extra operational cost for sludge disposal.
Sludge production, extra operational cost for sludge disposal.
Subsequent treatments are required to improve the removal efficiency of heavy metal
Not all ion exchange resin is suitable for metal removal, high capital cost
High operational cost, prone to membrane fouling
Costly, prone to membrane fouling
High energy consumption due to high pressure required (20–100bar), susceptible to membrane fouling
Low selectivity, production of waste products
[86-89]
[87]
[87, 90]
[91, 92]
[91, 93]
[87]
[87, 91]
[81, 94, 95]
. Flotation presents several advantages over the more conventional methods, such as high removal efficiency, high selectivity of metal ions, low detention periods, low operating cost production of more concentrated sludge and high overflow rates [96]. Operation costs, high initial capital cost and high maintenance are the disadvantages of this process.
Electrochemical heavy metal wastewater treatment technologies are considered as rapid and well-controlled that require fewer chemicals, offer good reduction yields and generate less sludge. On the other hand, electrochemical methods involving high cost electricity supply and high initial capital investment, this restricts the technique development.
Biological technologies by using different low materials were found be very effective techniques with higher uptake percentage. Although biological techniques are low cost and friendly methods for the environment they require large areas and proper operation and maintenance [40].
Even though all above techniques can be employed for the remediation of heavy metal wastewater, it is important to remarkable that the selection of the most suitable remidation methods depends on the initial concentration of the metal, wastewater component, plant flexibility and accuracy, capital investment, operational cost and environmental impact, ...etc [60, 85].
1.4. Enviromental pollution with iron metal and its removal:
Iron is considered the second metal among the most abundant metals on the earth crust. In the periodic table of elements, iron occupies the 26th elemental position. Iron is existing in many forms in water as shown in Figure (6) [97]. Biologically it is a most crucial element for survival and growth of almost all living organisms. As it is the cofactor for many vital enzymes and proteins. It is one of the vital constituents of organisms like algae and of enzymes such as catalase and cytochromes, in addition to oxygen transporting proteins, such as myoglobin and hemoglobin. Because of iron inter-conversion between ferrous (Fe2+) and ferric (Fe3+) ions, it is regarded as an attractive transition metal for various biological redox processes owing. The iron source in surface water is anthropogenic and is associated with mining activities. Sulphuric acid production and the discharge of ferrous (Fe2+) occurs because of iron pyrites (FeS2) oxidation that are common in coal seams. [98-100]. The following equations represent the simplified oxidation reaction for ferrous and ferric iron [99]:
2FeS2 + 7O2 + 2H2O 2FeSO4 + 2H2SO4 (ferrous)
4FeSO4 + O2 + 10H2O 4Fe(OH)3 + 4H2SO4 (ferric)
Mediated reactions of iron support the respiration process of most of the aerobic organisms. If it is not shielded correctly, it can catalyze the reactions involving radicals formation which can destroy biomolecules, tissues, cells and the whole organism. Iron poisoning has always been a subject of interest chiefly to pediatricians. Children are highly susceptible to iron toxicity as they are exposed to a maximum of products containing iron [101].
Figure (6): Classification of different forms of Iron presen in water.
Iron toxicosis occurs in four stages [100]:
The first stage which takes place after 6 hrs of iron overdose is noticeable by effects of gastrointestinal such as vomiting, diarrhea and gastro intestinal bleeding.
The second stage progresses within (6 - 24hrs) of overdose and it is regarded as the inherent period, a period of apparent medical recovery.
The third stage happens between 12 to 96 hrs after certain clinical symptoms onset. This stage is characterized by shocks, tachycardia, lethargy, hypotension, metabolic acidosis hepatic necrosis, and sometimes death.
The fourth stage take place in betwwen 2 to 6 weeks of iron overdose. This stage is marked by the gastrointestinal ulcerations formation and strictures development.
Iron uptake excess is a serious problem in meat eating and developed countries and it increases the cancer risk. Workers who are highly susceptible to asbestos that contains almost 30% of iron are at high risk of asbestosis, which is the second most important reason for lung cancer. It is said that asbestos associated cancer is related to free radicals. Loose intracellular iron can also promote DNA destruction. Iron can initiate cancer mainly by the DNA oxidation process.
Iron salts such as iron sulfate, iron sulfate heptahydrate and iron sulfate monohydrate are of low acute toxicity when exposure is through dermal, oral and inhalation routes and hence they have been placed in toxicity category 3. Moreover, the Food and Drug Administration considered that iron salts are safe and their toxic effects are very much negligible.
Free radicals formation is the outcome of the toxicity of iron. During pathological and normal cell processing, byproducts such as hydrogen peroxide and superoxide are produced, which are regareded to be free radicals. These free radicals are actually neutralized by enzymes such as catalase superoxide dismutase, and glutathione peroxidase but the superoxide molecule has the capability to release iron from ferritin and that free iron reacts with more and more of hydrogen peroxide and superoxide forming free radicals with high toxicity such as hydroxyl radical. Hydroxyl radicals are dangerous as they can initiate lipid peroxidation, inactivate certain enzymes, cause DNA strand breaks and can depolymerize polysaccharides. This can occasionally lead to cell death.[100, 102, 103]
Tahir and Rauf [104] studied the Removal of Fe(II) from the galvanized pipe manufacturing industry wastewater by adsorption onto bentonite clay. The adsorption of Fe(II) from aqueous solutions over a concentration range from 80 to 200 mg/l, shaking time of 1–60 min, adsorbent dosage 0.02 – 2 g and pH of 3. The process of removal follows both the Langmuir and Freundlich isotherm models and also obey the first-order kinetics. The maximum removal (> 98%) was observed at pH of 3, 0.5 g of bentonite with initial concentration of 100 mg/l. The Fe(II) removal efficiency was also tested using wastewater from a galvanized pipe manufacturing industry. Higher than 90% of Fe(II) can be effectively removed from the wastewater by using 2.0 g of the bentonite. The effect of cations (i.e. manganese, cadmium, lead, chromium, zinc, copper, nickel and cobalt) on the removal of Fe(II) was studied in the concentration range of 10–500 mg/l. All the added cations reduced the Fe(II) adsorption at high concentrations except Zn. Column studies have also been investigated using a certain concentration of wastewater. More than 99% recovery has been attained by using 5 g of the bentonite with nitric acid solution (3 M).
Das et al. [105] in a study investigated that the traditional method of using ash for disposal of iron from groundwater can eliminate iron to desired level without increasing the pH behind the acceptable limit. The banana pseudostem ash is among the different plant ashes used for iron removal. It has been found to be most appropriate for iron elemination. The ash improves iron uptake. The designed iron elemination system is expected to be convenient for household use. The optimum values of the different parameters for iron removal are 200–300 mg L−1 ash, 1.0 L h−1 rate of filtration and time of residence (1h) for groundwater having 2.20 ppm iron concentration. For groundwater having higher [Fe], the amount of ash can be increased and can be decreased gradually throughout continuous use. The technique has the advantages of low manufacturing cost, almost nil recurring cost, viz., simplicity in use, no electricity requirement, and increasing the essential minerals such as K, Ca in the treated water.
Al-Anber et al. [106] in their study the batch removal of Fe3+ from aqueous model solution under different experimental conditions using Jordanian natural zeolite (JNZ) has been investigated. The contact time influences, initial concentration of metal, temperature and concentration of adsorbent dosage have been studied. The adsorption efficiencies are found to be residence time dependent, increasing the contact time in the range between 1 and 150 min. The sorption equilibrium has achieved between 60 and 150 min. The optimum adsorption has occur at 30°C of temperature. The equilibrium adsorption capacity of JNZ adsorbent used for Fe3+ were evaluated and extrapolated using Freundlich and Langmuir isotherm models and the experimental data are found to fit Langmuir isotherm more than Freundlich isotherm.
Ghosh and his team [107] reported the results of astudy on electrocoagulation (with electrodes from aluminum) for iron Fe(II) elemination from aqueous medium. The removal of Fe(II) was composed of two principal steps; (a) oxidation of Fe(II) to Fe(III) and (b) subsequent Fe(III) disposal by the freshly formed aluminum hydroxides complexes by adsorption/surface complexation followed by precipitation. Experiments were executed with various current densities ranging between 0.01 and 0.04 A/m2. Other parameters such as salt concentration, pH and conductivity were maintaned constant as per tap water quality. It was observed that as the current densities increase, the elemination of Fe(II) increased. Satisfactory iron removal of around 99.2% was attained at the end of 35 min of operation from 25 ppm initial Fe(II) concentration.
Vasudevan et al [108] reported the results of astudy on the uptake of Iron from drinking water by electrocoagulation using galvanized iron as the cathode and magnesium as the anode. Experiments were done as a function of current density, pH and temperature. By using both the Langmuir and the Freundlich isotherm models, the adsorption capacity was estimated. The results demonstrated that the maximum efficiency of removal of 98.4% was obtained at 0.06 A dm– 2 of current density and pH of 6.0. The adsorption of iron was better illustrated by fitting the Langmuir adsorption isotherm, which suggests adsorbed molecules monolayer coverage. The adsorption process followed a kinetics model of second-order. Temperature studies illustrate that adsorption was endothermic and spontaneous in nature.
Bulai and Cioanca [109] in astudy found that Purolite S930 is an effective sorbent for Iron (II) ions uptake from aqueous model solutions in different operating conditions. The Iron (II) elemination percent has a maximum at pH 5.0 and increases with contact time and resin dose increasing and decreases with solution initial concentration increasing.
Wang [110] investigated on crushed concrete and limestone removed Fe(II) from synthetic groundwater in laboratory columns. The results shown that achieving average Fe(II) uptake of greater than 216 (approximately 133 L treated) and 99% over 288 (approximately 172 L treated) pore volumes, for crushed concrete and limestone, respectively. Calcium siderite which formed in limestone columns as a form of precipitate; this formation had no significant effect on porosity of the system, but may have impeded Fe(II) remidation by limiting available surface area for adsorption. Results suggested that field-scale passive iron removal systems, using similar materials,merit exploration. Because differences are expected, pilot-scalefield tests are warranted.
Nandeshwar et al. [111] in astudy foucsed on iron uptake from real wastewater samples of Nag River, India using Green activated carbons from different waste materials such as orange peels, sawdust, C. procera leaves and coconut shells. All the selected waste materials were carbonized in muffle furnace and activated using various agents such as HCl, HNO3, and H2SO4. The results showed that all adsorbents have the potential capacity to separate iron, which further highly increases after its activation. The most promising green adsorbents were found to be orange peels and HCl was the best activating agent. The order of iron uptake from wastewater is: orange peels then coconut shells then sawdust then C. procera leaves. Similarly it was found that charcoal activated with HCl can separate around 77–90% iron followed by HNO3 (70–80%) and H2SO4 (58–75%).
Vries et al. [112] conducted astudy on iron separation from water by using rapid sand filtration. A model has been developed that takes into account the main properites of (submerged) rapid filtration: the water quality parameters of the influent water, marked pH, concentration of iron(II), homogeneous oxidation in the supernatant layer, surface sorption and heterogeneous oxidation kinetics in the filter, and adsorption characteristics of the filter media. Adsorption isotherm data collected from different Dutch remidation places show that Fe(II) adsorption may vary strongly between them, but generally increases with higher pH. The model has a sensitivity for (experimentally) determination of adsorption parameters and the heterogeneous oxidation rate.
Wang et al. [113] in astudy focused on Effects of solution chemistry on the removal reaction between Fe(II) and calcium carbonate-based materials by using a permeable reactive barrier consist of calcium carbonate-based materials (CCBMs), such as limestone. There is no significant effect on the uptake of Fe(II) by limestone from pH 7 to 9. Na+ significantly affected elemination of Fe(II) at levels of 100 mg/L and above. Ca2+ and Mn2+ showed effect on removal as low as 10 ppm Ca2+ and 5 ppm Mn2+. natural organic matter (NOM) premixed with Fe(II) (10 ppm Dissolved organic carbon (DOC) ) resulted in final Fe(II) levels above GCTL (groundwater cleanup target level). NOM retained 0.05 mg Fe(II)/mg for 2/3 sources and 0.032 mg/mg for 1/3.
Indah and Helard [114] conducted astudy on Evaluation of Iron and Manganese-coated Pumice from Sungai Pasak, West Sumatera, Indonesia for Fe (II) and Mn (II) Removal from aqueous model solutions. The effect of soaking time for iron and manganese coating was studied and as comparison. The experiments were performed in batch mode at room temperature between 20 and 25 oC, pH 7; adsorbent dose of 10 g/L; adsorbent diameters of 0.30-0.50 mm; 90 minutes of soaking time and100 rpm of agitation speed. The results showed that the optimum soaking time for manganese coating and iron for removal of Fe (II) and Mn (II) was 100 hours. Iron-coated pumice showed to have high removal efficiency compared to uncoated and manganese-coated pumice. More than 84% of Fe(II) with 15 ppm initial concentration was removed by 10 g/L iron-coated pumice, while by using uncoated and manganese-coated pumice, the elemination efficiencies were less than 75% . The desorption study noticed that up to 20% of Fe (II) was recovered from the three kinds of pumice adsorbent. Overall research indicated that pumice from Sungai Pasak may be a promising adsorbent for iron disposal from water and wastewater.
1.5 Enviromental pollution with lead metal and its removal:
Lead is a highly toxic metal whose widespread use has give rise to comprehensive environmental pollution and problems of health in many world parts. Lead is a bright silvery metal, slightly bluish in a dry atmosphere. It begins to tarnish with air contact, thereby forming a complex mixture of compounds, depending on the given conditions. Figure (7) shows various sources of lead pollution in the environment [115].
The lead exposure sources involve mainly industrial processes, smoking and food, drinking water and domestic water sources. The lead sources were gasoline and house paint, which has been extended to plumbing pipes, lead bullets, storage batteries, pewter pitchers, faucets and toys [116]. larger than 100 to 200,000 tons of lead per year is being emitted from car exhausts in the US. Some is taken up by plants, fixation to soil and flow into water bodies, hence human exposure of lead in the general population is either owing to drinking water or food. Lead is an extremely toxic heavy metal which disturbs different physiological processes of plant and unlike other metals, such as, copper, manganese and zinc, it does not play any biological functions. A plant with high concentration of lead fastens the reactive oxygen species (ROS) production, resulting in damage of lipid membrane that finally leads to destruction of photosynthetic and chlorophyll processes and suppresses the plant overall growth [117]. Some research stated that lead is capable of suppressing the tea plant growth by reducing biomass and debases the tea quality by changing the quality of its components [118]. Even at low concentrations, lead remediation was found to cause large instability in ion uptake by plants, which in turn leads to significant metabolic changes in photosynthetic capacity and ultimately in a strong inhibition of plant growth.
Figure (7): Various sources of lead pollution in the environment.
Poisoning of Lead was considered to be a classic disease and the marks that were seen in children and adults were fundamentally attached to the gastrointestinal tract and the central nervous system [119]. Lead poisoning can also take place from drinking water. The pipes which carry the water may be made of lead and its compounds which can pollute the water [120]. According to the Environmental Protection Agency (EPA), lead is regarded a carcinogen. Lead has large effects on different parts of the body. Distribution of Lead in the body initially based on the blood flow into various tissues and almost 95% of lead is precipitated in the form of insoluble phosphate in skeletal bones [121]. Toxicity of lead, also called lead poisoning, can be either chronic or acute. Acute exposure can result in headache, abdominal pain, appetite loss, renal dysfunction, vertigo, sleeplessness, arthritis, hallucinations, hypertension and fatigue. Acute exposure chiefly occurs in the work place and in some manufacturing industries which make use of lead. chronic exposure of lead can cause psychosis, autism, weight loss, allergies, mental retardation, dyslexia, hyperactivity, kidney damage, paralysis, muscular weakness, birth defects, brain damage, and may even lead to death [122].
Eventhough lead toxicity is preventable it still remains a dangerous disease which has effect on most of the organs. The plasma membrane moves into the brain interstitial spaces when the blood brain barrier is exposed to great levels of lead concentration, leading a condition called edema. It disrupts the intracellular second messenger systems and alters the the central nervous system functioning, whose protection is highly important. Domestic and environmental lead ions sources are the main reason of the disease but with appropriate precautionary measures it is possible to reduce the risk correlated with lead toxicity [120].
Generally, Impact of lead exposure in humans has been known to cause wide variety of health problems such as [123] :
• Various forms of blood disorders and Anemia
• Rapid deterioration of brain and the nervous system
• fertility decreasing both in men and women
• Failure of the kidney
• Alzheimer disease
Many studies have been reported for lead elemination from aqueous solutions. Pala and Dursun [124] studied that the results of a study on adsorption of Pb (II) ions from artificial contaminated tap water by using a natural zeolite (Clinoptilolite). Clinoptilolite mineral which has mesh size of 25-140 was used by activating with HCl, and the efficiencies of lead ion disposal were evaluated. Experiments were occured under laboratory batch conditions were run at different values of pH, temperatures. The highest efficiency of removal was found as about 87% at pH 5. In similar way, experiments were done at different temperature values, and the utmos efficiency was achieved at 30oC. The efficiency obtained under these conditions was 89.95%. The highest lead disposal efficiency was achieved with shaking speed of 200 rpm.
Mavropoulos et al. [125] in their study found that the composite of hydroxyapatite-alginate was effective in the elemination of lead ions and lead phosphate nanoparticles from high-polluted simulated gastric fluid. The cross-linked polymer chain had a double role: (i) keep Pb2+ ions and lead phosphate nanoparticles bounded to the surface of bead, impeding their bioavailability in stomach fluid; and (ii) delay dissolution of HA in the stomach acidic conditions, confirming that an excess of Ca2+ will not be released to simulated gastric fluid. Desorption studies in simulated enteric fluid stated that lead stayed immobilized in the calcium phosphate phase in the intestinal tract. These results indicate HA–alginate composite as effective system for heavy metals disposal from polluted gastric and enteric human fluids, reducing its adsorption by the human body.
Meski and his team [126] showed in their work with hydroxyapatite prepared from the egg Shell , that the carbonate hydroxyapatite prepared from egg shell (CHAPF) represents the highest capacity for Pb2+ ions adsorption from aqueous solution. It has been found that the initial adsorption rate was high. The sorption process obey the model of Langmuir isotherm with low temperature dependency and high adsorption capacities. The thermodynamic functions were calculated, and it can be concluded that the Pb2+ adsorption over CHAPF is an exothermic and spontaneous process. The adsorption was greatly pH dependent, with a high uptake of lead at pH = 3. These results show that the lead uptake by CHAPF was very sensitive to the initial concentration of Pb2+ in aqueous solution. For the high concentrations [(500 to 700) mg·L-1], two stages were observed: Pb2+ ions adsorption on the CHAPF surface and an ion exchange reaction between Ca2+ of CHAPF and Pb2+ ions in aqueous model solution.
Shrestha et al. [127] conducted a study on lead (II) disposal from aqueous solutions using prepared activated carbon. Two series of carbon have been synthesized from Lapsi seed stones by treating with concentrated H2SO4 and HNO3 in a mixture with H2SO4 in the ratio of 1:1 by weight for disposal of metal ions. pH 5 was the optimum pH for lead adsorption. For the equilibrium isotherms description, the adsorption data were better fitted with the Langmuir adsorption equations than Freundlich equation. The maximum adsorption capacity of Pb (II) on the produced activated carbons was 277.8 mg/g with a mixture of HNO3 and H2SO4 and 423.7 mg/g with H2SO4. The waste material used in activated carbons preparation is readily available and cheap. Therefore the carbons synthesized from Lapsi seed stones can work as potential low cost adsorbents for the elemination of Pb (II) from water.
Jalali [128] investigated astudy on stalk of Sunflower, an agricultural waste, acts as an adsorbent for the cadmium and lead disposal from aqueous solutions. Adsorbent was synthesized by washing residue of sunflower with deionized water until the solution become colorless. The results stated that the adsorbent has good sorption potential and maximum removal of metal was detected at pH 5. Within 150 min of operation about 97 of Pb ions were eleminated from the effluents. Curves of lead sorption were well fitted to the modified two-site Langmuir model. Lead adsorption capacities at optimum operation conditions were 182 mg/g. The kinetics of Pb ions adsorption from aqueous model solutions were also analyzed. The experimental data were foud to be fitted to pseudo-second-order kinetic model. fitted models (R2 > 0.999). The maximum adsorption capacity for Pb(II) ions adsorbed onto entrapped silica nanopowders was evaluated to be 83.33 mg/g.
Soltani and his team[129] conducted astudy on adsorption of Pb(II) ions from the aqueous solution by using entrapped silica nanopowders within calcium alginate in order that determination the thermodynamic, isotherm and kinetic of the adsorption process. According to the results, an initial pH of 5.0 was found to be optimal for the Pb(II) ions adsorption. The capacity of adsorption reached to 36.51 mg/g with increasing the contact time to 180 min at 50 ppm as initial Pb(II) ions concentration. However, the equilibrium contact was estimated to be 90 min owing to no significant increase in adsorption effeciency after this time. The results of studies stated that the isotherm of Langmuir and pseudo-second order model of kinetic were the best.
Yarkandi (2014) [130] carried out batch experiments for lead separation from waste water using natural american bentonite and activated carbon. The results show that the amount of Pb++ adsorption increases with solution pH, initial concentration of metal ion and contact time but decreases with temperatures and amount of adsorbent. The adsorption process has well fit pseudo-second order kinetic model. Langmuir and Freundich adsorption isotherm models were found to be applicable to the adsorption process where both were applies to analyze adsorption data. Thermodynamic parameters e.g. ΔH°, ΔG° and ΔS° of the adsorption process was found to be endothermic. Finally it can be seen that activated carbon was found to be less effective for disposal of Pb+2 ions than bentonite.
Bartczak et al [131] in a study investigating the lead (II) ions adsorption from aqueous model solution on peat as adsorbent with low-cost, observed that The sorption capacities of peat with respect to lead(II) ions was 82.31 mg(Pb2+)/g. Slightly well results were achieved with adsorption efficiency reached 100% just after 3 min (15 and 30 ppm), 5 min (50 ppm) or 15 min (100 ppm) of the process. It indicates the utmost affinity of the peat surface for lead ions. the optimum adsorbent mass was found to be 5 g/L. To prevent metal precipitation as hydroxides and including the obtained results pH = 5 confirmed as an ideal.
Sangeetha et al [132] investigated a study on lead ions disposal from aqueous solution by using novel hydroxyapatite/alginate/gelatin composites. Pb2+ elemination ability of wet precipitation synthesized biosorbents HA/Alg and HA/Alg/Gel has been investigated for different dosages. Complete disposal was obtained from 7 to 24 hours by both the adsorbents. The sorption kinetics were found to be best fit to the pseudo-second order equation and the equilibrium well followed Langmuir isotherm model. The elemination capacity was higher for lower dosage studied and the rate of disposal was higher for the higher dosage studied.The sorption mechanism involved was dissolution/precipitation, ion exchange and surface complexation process. according to The results, it can be seen that both the composites under study are potential candidates for Pb2+ disposal and precisely gelatin enhanced the maximum sorption than the alginate alone which is composed with the hydroxyapatite.
Cheraghi et al. [133] showed in their work with waste tea leaves, upon parameters optimization like initial metal concentration, Temperature, pH and adsorbent amount, that maximum uptake efficiency was achieved at pH 6. Also, as the initial metal concentration decreased the adsorption of Pb (II) ions increased. the equilibrium adsorption isotherm data fits well with the Langmuir isotherm model and its calculated maximum adsorption capacity of monolayer was 166.6 mg/g at 25±0.1˚C. The sorption kinetics were found to be best fit to the pseudo-second order equation.
Kanyal and Bhatt [134] reported the results of astudy on adsorption of Pb (II) ions from waste water by using household waste as an adsorbent. Banana peels, Pumpkins and Chicken eggshells are considered as good adsorbents for elemenation of heavy metals from polluted water. The effects of various parameters such as agitation speed, pH and residence time were examined and best results were observed at pH 7, 90 mins and 100 rpm. The results stated that household waste usage such as these can be act as a good biosorbent for disposal of heavy metals on a large scale and establish effective, and inexpensive techniques in wastewater remidation.
Wang et al. [135] carried out a study for selective disposal of lead and other metals such as cadmium and copper from wastewater by gelation with alginate for effective recovery of metal. The results evaluated that gels can be formed speedily between the metals and alginate in lower than 10 min and the rates of gelation fit well with the pseudo second-order kinetic model. The optimum ratio of dosing of alginate to the metal ions was found to be between 2:1 and 3:1 for Pb2+ elemination and around 4:1 for Cu2+ and Cd2+ elemination from wastewater, and the metal disposal efficiency by gelation increased with the solution pH. Alginate has a higher affinity of gelation toward Pb2+ than Cd2+ and Cu2+, which permitted a selective uptake of Pb2+ from the wastewater in the existing of Cd2+ and Cu2+ ions.
Liu et al. [136] focused on the disposal of lead ion from waste water using hydroxyapatite scaffolds synthesized from scales of fish. Powder of fish scale obtained from Tilapia fish (Oreochromis mossambicus) was used for preparing scaffolds for lead removal. The maximum adsorption capacities (qmax) were 344.8 mg/g and 208.3 mg/g in solutions pH of 2.2 and 5, respectively. More than 99.9% of the lead ion was eleminated after 20 min.
Al Lafi et al. [137] conducted astudy on Lead removal from aqueous aolutions by polyethylene waste/nano-manganese dioxide composite . The adsorption results investigated that the synthesized adsorbent can effectively eleminate Pb+2 ions from aqueous solutions, with a maximum adsorption capacity of 50.5 mg/g. Almost 60 % of the initial Pb+2 concentration were adsorbed within the first hour, and it was concluded that 2 h was the optimum time for Pb+2 elemination . A pH value of 5.0 was determined as an optimum and was used for the rest of this study. Regeneration of the composite can be performed using 0.5 mol/L HCl solution with Pb+2 percentage of recovery reached 95 %. It also efficiently adsorbed Pb2+ after five sorption/desorption cycles with 84 % as percentage of removal.
Heraldy [138] conducted a study of sorption of Pb(II) from the aqueous mediums using biosorbent synthesized from waste of tomato and residue of apple juice (AR). The optimum conditions for maximum removal percentage of Pb(II) by biosorbents were found to be 0.1 g of sorbent at pH 4.0 and 90 min contact time for tomato waste and 60 min for AR. The experimental data were found to be well matched with Freundlich than Langmuir isotherm model. A kinetic study showed that Pb(II) sorption follows the pseudo-second-order kinetics, which confirms that AR and waste of tomato biosorbents are 108 and 152 mg/g, respectively.
1.6 Role of alginate in heavy metals removal:
Alginate is belong to the anionic polymers family which is naturally occurring usually produced from brown seaweed, and has been widely investigated and used for many biomedical applications, due to its relatively low cost, low toxicity, biocompatibility, and mild gelation by divalent cations addition such as Ca2+ [139]. Alginates are a polysaccharides composed of variable ratios of β-D-mannuronate (M) and its C-5 epimer α-Lguluronate (G) linked by 1–4 glycosidic bonds (Fig.8). In the 1880s, alginates were first separated from brown seaweeds, and its production for commerce started in the early 20th century.
The production of alginate can be carried out by two bacteria genera, Azotobacter and Pseudomonas and various genera of brown seaweed and [140]. Alginate which is available in commerce can usually extracted from brown algae (Phaeophyceae), including Laminaria digitata, Macrocystis pyrifera, Laminaria japonica, Ascophyllum nodosum and Laminaria hyperborea by remidation with aqueous alkali solutions, usually with NaOH. The extract is filtered, and in order to precipitate alginate either calcium or sodium chloride is added to the filtrate. By treatment with dilute HCl, the alginate salt can be converted to alginic acid. After further purification and conversion, power of water-soluble sodium alginate is generated. [139, 141]
Figure (8): Chemical structure of alginate. M – mannuronate residues, G – guluronate residues.
The synthesised alginates were found to be extremely effective in uptake of heavy metal ions from aqueous model solutions. Removal of more than 93% Cr(VI) ions was obtained from aqueous solution in batch process using this type of biosorbent. [142, 143]
Alginate is regarded a biopolymer with many applications in food industry, drug delivery systems, cosmetics and cell encapsulation. In wastewater remediation could play a significant role in disposal heavy metal ions due its advantages, such as biodegradability, facile obtaining procedure, economical, biocompatibility and environmental friendly [144].
NiŃa et al. [144] in astudy found that calcium alginate microparticles has a good affinity for the divalent cations. The heavy metals adsorption was examined as a function of residence time between the samples of synthetic wastewater and the alginate and the polymeric beads morphology. The microparticles of calcium alginate were synthesized using a laboratory procedure, sodium alginate aqueous solution was dropped in a solution of calcium chloride. The polymeric microparticles which have a controlled porosity seem to be an appropiate alternative to develop a aprocess for elemination of heavy metal from industrial wastewater.
Singh et al. [145] investigated astudy on effective removal of Cu2+ ions from aqueous medium using alginate as biosorbent. Maximum removal of Cu2+ ions (85.3%) from aqueous medium was observed at pH 5.5, alginate dosage of 2.5% and initial copper concentration of 275 ppm with 50 min as agitation time. Thus, the reultant experimental data has been fitted well with both Langmuir and Freundlich isotherm models.
1.7. Heavy metals removal by bentonite and kaolin clays:
Clay is one of potential good adsorbent substitutes to other adsorbent owing to its layered structure, large surface area, mechanical and chemical stability and high capacity of cation exchange. The existence of two acidity types, Lewis and bronsted in clays increases the clays adsorption capacity. Aluminum oxides and clay minerals such as bentonite and kaolin, are the most wide-spread minerals of the earth crust which are known to be good adsorbent of different metal ions, organic ligands and inorganic anions [146, 147].
1.7.1. Bentonite:
Bentonite is considered a strong candidate as an adsorbent for heavy metal elimination due to its abundance and its low cost. Bentonite as a representative clay mineral is a clay chiefly consist of montmorillonite, a 2:1 type of aluminosilicate Bentonites are extremely valued for their sorption [146, 147].
Zia and his team [148], evaluated the removal of Pb2+ and Cd2+ by using Methionine modified bentonite/Alginate (Meth-bent/Alg) nano composite. The desorption study presents that 99% of the adsorbed Pb2+ and Cd2+ can be desorbed by using oxalic acid (0.1 M) as eluting agent with regeneration ability up to fifth cycle effectively.
Wu et al. [149] reported the adsorption of Th4+ from aqueous solution by using Novel magnetic organo-bentonite-Fe3O4 (polysodium acrylate) (OB-Fe3O4 PSA) superabsorbent nanocomposites. OB-Fe3O4 PSA super absorbent could be regenerated through the desorption of Th4+ using HCl solution (0.1 mol/L) and the adsorption capacity was still greater than 3.6 mmol/g after five successive adsorption–desorption processes.
Tan and Ting [150] reported that plain alginate and alginate immobilized bentonite beads have good reusability potential. remidation with HCl (10 mM) successfully eluted 93.05% and 94.33% of the Cu2+ ions loaded onto plain alginate and alginate immobilized bentonite, respectively, after three cycles of sorption–desorption test. There was no significant difference in the percentage of Cu2+ desorped in the three sorption–desorption cycles for both the plain alginate (93.15%, 93.54% and 92.48%) and immobilized-bentonite (92.38%, 96.03% and 94.75%). This established high reusability of the developed immobilized bentonite without remarkable losses in their Cu2+ disposal capacities.
1.7.2. Kaolin:
Kaolin is considered a type of clay rock, which incloses some chemical elements such as Na, Al, Mg, Ti, Ca, Fe, Ka, Si , and so on. The silicon mass proportion is more than 50%. Kaolin can be divided into two classes, which is coal-series kaolin and nature kaolinite The monolayer crystal structure of kaolin is comprised of siliconoxygen tetrahedral sheets and aluminum-oxygen octahedral sheets, just like the molecular structure model of kaolin shown in Figure (9) [151]. Metal ions removal using kaolinite clay is depend on mechanisms of ion exchange and adsorption and kaolinite has a relative low capacity of cation-exchange (CEC) [3–15 meq/100 g of clay] and smaller surface area ranged between 10 and 20 m2/g [152].
Figure (9): The molecular structure model of the kaolin.
Li et al. [153] reported that a novel environmental friendly material, calcium alginate immobilized kaolin (kaolin/CA), which synthesized using a sol-gel method, have good effeciency for copper uptake from waste water. the experimental adsorption was described using the Langmuir isotherm, the maximum capacity of Cu2+ adsorption by the kaolin/CA reached up to 53.63 mg/g. The thermodynamic studies indicated that the adsorption reaction was found to be an endothermic and spontaneous process.
Yavuz and his team [154] investigated the elemination of heavy metals such as Cu(II), Ni(II), Co(II) and Mn(II) from aqueous solution using raw kaolinite. The sorption of these metals on kaolinite conformed to Langmuir adsorption equation. Langmuir Cm constants for each metal were found as 0.919 mg/g (Co), 10.787 mg/g (Cu), 1.669 mg/g (Ni), 0.446 mg/g (Mn), at 25 oC, respectively. Also, kinetic and thermodynamic parameters like entropy (ΔS), enthalpy (ΔH) and free energy (ΔG) were evaluted and indicated that heavy metal adsorption on kaolinite was an endothermic process and the process of adsorption was preferable at elevated temperatures.
Larakeb (2017) [155] evaluted the Zinc Removal from Water by Adsorpion on Kaolin and Bentonite clays. The kinetics of adsorption results showed that zinc disposal is max. with and 45.48℅ efficiency for kaolin after 60 min of residence time and after 20 min with 89.8 ℅ efficiency for bentonite. Adsorbent dose increasing from 0.5 to 8 g/l enhance zinc elemination efficiency for 5 ppm like an initial concentration. Zinc disposal efficiency by the two adsorbent decreases with rising of the initial Zn concentration from 2 to 20 ppm. pH of treatment has considerable effects on the retention rate of zinc. The efficiencies of Zn removal are noticeable at basic pH. Whatsoever reaction parameter tested, it appears that kaolin is less effective than bentonite.
1.8. Removal of heavy metals by gelatin:
It has been reported that gelatin has agood affinity for heav metals removal, individaly or combined in composites with other materials. Itabashi et al. [156] found that The good effect of copper elimination by gelatin was achieved by the foam treatment of this solution. And also lead was successfully eliminated by the same treatment. gelatin powder was used for the adsorption treatment to raise this effect before the foam treatment. About 99 % of copper in the range of pH between 6.5 and 7.3 and approximately 100% of lead at pH 7.0 was eliminated respectively. Hayeeye et al [157] in their study found that, The maximum capacity of adsorption of gelatin/ activated carbon for Pb2+ ions was obtained to be 370.37 mg g-1. The separation process for Pb2+ ions was found to be relatively rapid with 92.15% of the adsorption finished in about 5 min as residence time in batch conditions. Adsorption was achieved at pH value as low as 2.0 and maximum adsorption was observed at a pH of almost 5.
1.9. Hydroxy apatite from bio waste materials:
Hydroxyapatite (HAP, Ca10(PO4)6(OH)2), a naturally available form of calcium phosphate and a component of hard tissues, has been reported to work as an efficient ion uptake material for different heavy metals from aqueous medium owing to its low solubility of water and excellent reactivity. The high stability of HAP structure, along with its flexibility permit a high variety of exchanges (particularly Ca ions with divalent heavy metal ions, such as As, Cd, Cu, Zn, Pb, Co, Ni, Sb,U, Hg, of huge importance in the environmental science field [158-160]. HAP can be extracted from different biowaste such as, eggshells and bovine bones. It can be synthesized through various methods which can be generally divided into two major routes: solid state reaction and wet methods [161]. including sol-gel technique, wet precipitation, hydrothermal process, mechanochemical method. Depending on the used techniques, HAP with several morphologies, composition, specific surface and crystalline degree have been obtained and appear to have different effects on the mechanical properties, bioactivity and dissolution behavior in biological environment [160, 162].
1.9.1. Preparation of hydroxyl apatite by sol-gel technique:
Sol-gel method is used for obtaining HAP powder of fine particle (nanoparticle size). In the sol-gel method of the HAP the calcium compounds and phosphorus precursors are transformed through condensation and hydrolysis reactions to the amorphous gels, which are further converted to ceramics when heated at comparatively low temperature. The polycondensation and hydrolysis are not separated in time, but occurs simultaneously [163]. Ceramic materials prepared by sol-gel route present many advantages over the others, such as homogeneous composition, low synthesis temperature and high product purity. Additional advantage of the sol-gel method is its applicability for surface coating.
There is no any form of secondary environmental damage as a result of high biocompatibility and its slightly- alkaline pH. The efficiency of HAP in eliminating heavy metal ions extremely depends on ion nature, diameter, charge and concentration, in addition to the treated water properties (temperature, pH) [160]
Putra and his team [164] showed in their work with eggshell, that for batch adsorption studies , at 90 min equilibrium time, 0.1 g biomass dosage and pH 6 were optimum biosorption conditions for Zinc and Copper ions elimination from aqueous mediums.
Agarwal and Gupta [165] in their study with eggshells focused mainly on evaluting varying concentrations (5, 10, 20, 40, 100 mg/L) of lead and copper ; this study reported a 92% - 100% removal of Cu when 0.5 to 1.5 g of eggshells (adsorbent) was used against 5 and 10 ppm of Cu; and adsorption efficiency of 80% to 100% for Pb at the same concentrations.
Deydier et al. [166] conducted a study for elemination of Pb from effluents using alow-cost material from meat and bone meal combustion residue. This residue was regarded as an apatite-rich material and was used as a low-cost substitute of hydroxyapatite in lead elemination from water . the mechanism was found to be as in pure apatite: surface complexation and dissolution of calcium hydroxyapatite, followed by lead hydroxyapatite precipitation.
Rohaizar et al. [167] in astudy, observed that pH = 7 and 350 rpm as an optimum agitation rate were ideal for copper elemination from water.
Avram et al. [160] in a study investigated that low crystallinity HAP which prepared by the direct reaction of diammonium hydrogen phosphate and calcium nitrate at alkaline pH, can be successfully used in heavy metal disposal from mine wastewater. For all the 10 metals studied (Zn, Cd, Pb, Mn, Co, Fe,Cr, Cu, Ni and Al), their content was fastly reduced by contact with HAP under the legal allowable limits for wastewater discharge in natural environment. The ion exchange importance in sorption processes was revealed and the pseudo-2nd order kinetics of manganese ions sorption on HAP was estimated.
1.10. Heavy metals removal by silica nanoparticles (Diatomeous):
Silica is used widely in nanoparticles coatings which used in water purification techniques. Silica coating activates the NPs surfaces having various functional groups owing to the abundant existing of silanol groups on the silica layer. It also prevents leaching low pH situations of NPs. It also facilitates the NPs with non- specific moieties, highly and group specific ligands. Polymer layered silicate nanocomposites possess improved properties at low filler contents. At neutral pH, as the particle size increase, the acidity of Si NPs will increase resulting in 5 to 20% ionisation of silanol groups, causing attraction between anionic Si surface and cations by ion pairing [168].
Surface-functionalized nanoporous silica, often referred to as self-assembled monolayers on mesoporous supports (SAMMS), has previously presented the ability to act as very effective sorbents for heavy metal uptake in a range of environmental and aquatic systems [169]. Diatomite is belong to the siliceous rock family, silicon dioxide is the main constituent of it with the proportions up to 90%. Diatomeous has some advantages such as wear resistance, heat resistance non-toxic and large specific surface area,..etc. Diatomite is a sort of polyporous material. The diatomite porosity is up to 90%, which means that diatomite has great adsorbability [151].
Soltani et al. [129] conducted astudy on adsorption of Pb(II) ions from the aqueous solution by using entrapped silica nanopowders within calcium alginate in order that determination the thermodynamic, isotherm and kinetic of the adsorption process. According to the results, an initial pH of 5.0 was found to be optimal for the Pb(II) ions adsorption. The capacity of adsorption reached to 36.51 mg/g with increasing the contact time to 180 min at 50 ppm as initial Pb(II) ions concentration. However, the equilibrium contact was estimated to be 90 min owing to no significant increase in adsorption effeciency after this time. The results of studies stated that the isotherm of Langmuir and pseudo-second order model of kinetic were the best.
Karnib and his team [170] used a composite from Activated Carbon, Silica and Silica Activated Carbon. Silica/AC (2:3) composite showed the greatest elemination percentage for 30 & 200 ppm nickel. SEM images revealed that AC was a microparticle with 25 μm as an average size, while silica were nanoparticles having an average size of 12 nm. Silica/AC (2:3) composite was the most effective microparticle for nickel disposal and it is highly recommended to be used in water treatment for its high adsorptive capacity followed by AC and silica nanoparticles.
Aim of the work
The presence of toxic heavy metals in water has caused several health problems with animals, plants, and human. So that the removal of toxic heavy metals from polluted waters are one of the most important issues of environmental remediation.
The development of new products which are abundant in nature, cheap and have no environmental impact for treatment of natural resources is an important area of material technology. Calcium alginate and its composites fulfills both characteristics and have the ability to eliminate heavy metals from industrial streams.
Hence the aim of the present work is to synthesize calcium alginate and different calcium alginate composites from clays and biowaste materials (egg shell and bovine bones) and their characterization using XRD, FTIR, EDX and SEM.
The second aim of this research is to use the prepared powder samples to remove two of most hazardous heavy metals (Pb2+ and Fe3+) and measure the efficiency of each sample for remediation process.
2. Experimental and Methods
2.1. Materials, Solutions and Chemicals:
Two different biowaste materials are used in preparation of Nano hydroxyapatite –calcium alginate composites are listed in Table (3).
Table (3): Raw biowaste materials and their sources.
Material Source
Egg shell Local Hen’s egg shells
Bovine bone Local Butcher shop (bovine femur bone)
High purity analytical grade chemical material have been used in the current study are listed in Table (4).
Table (4): Solutions and chemical materials used in the current study .
Materials Chemical composition Manufacture
Diammonium hydrogen phosphate (NH4)2HPO4 Oxford laboratory India
Ammonia solution NH4OH BDH Analar England
Lead nitrate Pb(NO3)2 BDH
Iron nitrate Fe(NO3)2.9H2O BDH
Calcium chloride CaCl2 ANALAR
NaOH solution NaOH WINLAP
HNO3 solution HNO3 WINLAP
Sodium alginate C6H7O6Na Oxford laboratory India
Gelatin - Oxford laboratory India
Raw material powders are prepared and their designation in the present study is given in Table (5).
Table (5): Designation of Raw materials source in the present study.
Raw material The Source
HAP (1) Egg shells calcined at 900 °C
HAP (2) Bovine bone calcined at 1000 °C
Bentonite Abu Zaabal Fertilizer & Chemicals Co.( originating from china )
Metakaoline Kaolin from Sinai Peninsula calcined at 800 °C
Diatomeous Kazakhstan
2.2 Preparation of calcium alginate composite powders:
2.2.1 Calcium alginate:
Calcium alginate powder is prepared by using controlled gellification method [171] with some modification reported by Daemi and Barikani [172]. CA Nanoparticles are obtained by addition of CaCl2 (0.05 M) to solution containing sodium alginate (3%) by mechanical stirrer at high stirring rates (Fig. 10). Six gram of polysaccharide is dissolved in 200 mL of deionized water with high rate stirring at room temperature for 1 h. After homogenization of sodium alginate solution by mechanical stirrer, the solution (1 litre) of 0.05 molar calcium chloride is added to the system. After 1 h of the rotation, it permitted to stand at room temperature for 24 hrs. prepared nanoparticles are purified by centrifugation for 30 min. The precipitate is washed and filtered three times using double distilled water to remove the adsorbed sodium and chloride ions. The filtered CA precipitate is dried at 60 °C for 12 hours in a dry oven. This dried solids is finely grinded and sieved below 63 μm before characterization and usage.
Figure (10): preparation of calcium alginate nanoparticles
2.2.2. Bentonite:
Bentonite of Abu Zaabal Fertilizer & Chemicals Company which originating from china (Table 5) is used as a raw material in the preparation of CA – Bentonite composite. Natural bentonite dried in the oven with a temperature of 80oC for 12 hrs with the aim to eliminate moisture, bentonite ground with mortal to break chunks of bentonite then calcinied at a temperature of 800oC for 2 hours which aims to eliminate Cl bond on bentonite, and the results in the form of nanoparticles of bentonite [173].
2.2.3. MetaKaolin:
Kaolin from Sinai Peninsula is used as a raw material in the current study. Kaolin is calcined at 800 oC for 2 hrs to obtain Metakaolin (Table 5) that used in preparation of CA – Metakaolin composite.
2.2.4. Diatomeous silica:
Diatomeous which originating from Kazakhastan (Table 5) is used as a raw material in the preparation of CA – Diatomeous composite. Diatomeous is dried at temperature of 80 oC for 12 hrs with the aim to eliminate moisture.
2.2.5. Hydroxy apatite:
HAP (I)
The Egg shells mainly contain calcium carbonate (91% - 94%), calcium phosphate (1%) and other organic matters, which makes it preferable for synthesizing CaO [174]. The Nano-hydroxyapatite HAP (I) is prepared from hen’s egg shells by a method described by Laonapakul [175] with little modifications. About 50 gm. of egg shells is boiled for 30 minutes in hot water and the protein membrane is removed manually. Egg shells are dried at 80 ºC for 6 hrs. Dried eggshells solid are grinded in the agate mortar into a fine powder. The fine eggshells powder is calcined at 900 ºC for two hours in order to remove any organic residue. At this temperature the eggshells convert into calcium oxide (CaO), according to the following reaction:
Ca CO3 + Heat → CaO + CO2↑
Calculating the stoichiometric of Ca / P molar ratio = 1: 0.67 solution is prepared from CaO and DAP. Firstly, about 8.4 gram of CaO is dissolved in 100 ml 2M HNO3 and then 11.885 gm of DAP is added dropwise to calcium solution while stirring and maintaining a stoichiometry of Ca/P ratio of 1.67. NH4OH solution is added dropwise to the mixture. The pH of the solution is maintained at pH=10. A white precipitate solution is obtained and vigorously stirred for 30 minutes and permitted to stand at room temperature for 24 hours. The left white precipitate is filtered and washed three times using double distilled water to remove the adsorbed ammonia and nitrate ions. In order to obtain the final HAP solid, the filtered hydroxyapatite precipitate cake is dried at 80 °C for 10 hours in a dry oven. This dried powder is heated at 700 °C for 2 hrs. in air using control electric muffle furnace, employing a heating rate of 10 °C/min. Afterwards, the calcined powder is finely grinded and sieved below 63 μm before characterization and usage.
HAP (II)
One of the raw materials in the present study is bovine femur bones obtained from local butcher market. HAP obtained from bovine bones is prepared by a method conducted by Agnieszka et al. [176] with little modifications .In the beginning, the bovine bones (about 150 gm) are crushed into small pieces (1-2 cm) and then boiled in 1 M NaOH for one hour then in hot water for 1.5 hrs. for defatting and easier removal of the organic residues and macroscopic adhering impurities. The bones are washed and cleaned well with water and for several times afterwards. The process is followed by drying the bones at 80°C for 6 hrs. to evaporate the adsorbed water. The solid bone pieces is calcined at 1000°C for 2 hrs. at heating rate 10°C /min. and afterwards are cooled slowly to room temperature. The final solid is grinded and sieved below 63μm, and kept for characterization and usage.
2.2.6. Calcium alginate - Composites (CACS):
The compsites are prepared by mixing the CA with approprate amounts of additaves (Bentonite, Metakaolin, Gelatin, HAP(1), HAP(2) and Diatomeous) with 2:1 ratio, respectively. 6 gram of polysaccharide (sodium alginate) is dissolved in 200 mL of deionized water with high rate stirring at room temperature for 1 hrs. After homogenization of sodium alginate solution by mechanical stirrer, three grams of the chosen additives is add, then 1 litre of 0.05M calcium chloride solution is added. After 1 hour of the rotation, it permitted to stand at room temperature for 24 hrs. The prepared nanoparticles are purified by centrifugation for 30 min. The precipitate is washed and filtered three times using double distilled water to remove the adsorbed sodium and chloride ions. The CA- composites precipitates are dried at 80 °C for 24 hours in a dry oven. This dried solid are finely grinded and sieved below 63 μm before characterization and usage. Table (6) shows The synthesized composites and its abbreviations.
Table (6): The synthesized composite:
Composite Compound Abbreviations
1 Calcium alginate-Bentonite CAB
2 Calcium alginate-Metakaolin CAMK
3 Calcium alginate-HAP(1) CAHA(I)
4 Calcium alginate-HAP(2) CAHA(II)
5 Calcium alginate-Diatomeous CAD
6 Calcium alginate-Gelatin CAG
2.3. Chemical analysis of raw materials:
A- Diatomeous Silica:
Table (7): Chemical composition of Diatomeous used in the present study (XRF fused bed)
Compound Wt. %
SiO2 71.50
Al2O3 10.40
CaO 0.74
Fe2O3 3.66
MgO 1.23
SO3-- 0.71
Na2O 0.78
K2O 1.25
Cl- 0.28
TiO2 1.04
P2O5 0.10
Mn2O3 0.02
Total 91.71
Loss on ignition 8.10
Total 99.80
B- Bentonite:
Table 8: Chemical composition of bentonite used in the present study (XRF fused bed)
Compound Wt. %
SiO2 55.11
Al2O3 17.27
CaO 0.99
Fe2O3 9.03
MgO 2.27
SO3-- 0.34
Na2O 3.67
K2O 1.19
Cl- 0.62
Total 90.49
Loss on ignition 9.42
Total 99.91
C- kaolin:
Table (9): Chemical composition of kaolin used in the present study (XRF fused bed)
Compound Wt. %
SiO2 47
Al2O3 37
CaO 0.20
Fe2O3 0.20
MgO 0.02
Na2O 0.15
K2O 0.04
TiO2 1.30
Total 85.91
Loss on ignition 13.40
Total 99.31
2.4. Preparation of heavy metal ion solutions:
The metal cations of lead (Pb2+) and Iron (Fe3+) are used in the present study in the form of salts: Pb(NO3)2 and Fe(NO3)3.9H2O. One liter stock solution of each metal cation is prepared using double distilled water with metal ion concentrations of 100 ppm.
2.5. Heavy metals uptake reaction:
The metal cations reactions are conducted as follow: 20 mg of each calcium alginate composite solids is equilibrated for different time periods (5, 10, 15, 20, 30, 60 and 120 min. respectively) in glass vials with 10 ml metal cation solution with continuous shaking. After different time intervals, the solid phases are separated by centrifugation, the supernatant solution was collected for chemical analysis using Perkin Elmer 2380 Atomic Absorption Spectro Photometer (Figure 11).
Fig. 11: Perkin Elmer 2380 Atomic Absorption SpectroPhotometer.
2.6. characterization of Composite solids:
The prepared Composite solids are characterized by using Scanning Electron Microscope (SEM), X-ray diffraction (XRD), Fourier Transform Infrared (FTIR) spectroscopy and Energy Dispersive Analysis X-ray (EDX) techniques.
2.6.1. X-ray diffraction (XRD):
An X-ray diffractometer (Philips X’ PERTMPD, America, with Cu Kα radiation, 40KV and 30mA) is used to determine the mineral phases and crystallinity of the different composite powders; (Figure12).
Fig. 12: Philips X-Ray Diffractometer
The specification criteria of XRD are adjusted at 2Ө range = 5° - 60° and λ = 1.54 Ǻ at a scanning speed of 2° /min. Each composite is used to fill the aluminum mold of the diffractometer with an average thickness of about 10 mm. The obtained phases are identified by correlation with the corresponding joint committee on powder diffraction standard card (JCPDS). The average crystallite size (D) of the obtained composite powders is calculated from XRD using the Scherrer formula [177] as shown below.
where:
λ= the wave length of the X-ray.
β = the full width at half maximum (FWHM) of the peak at the maximum intensity
Ө = the diffraction angle
2.6.2. Attenuated Total Reflection Fourier Transform Infrared (ATR-FTIR):
(ATR-FTIR) technique is used to determine the main constituent chemical functional groups of the different prepared samples and the type of chemical bonding between the different atoms existing in the groups. ATR-FTIR spectrometer (Bruker, Germany Alpha-p) is configured with ATR-FTIR sample cell including a diamond crystal with a scanning depth up to 2μm. Sample powders are applied to the surface of the crystal then locked in placed with a”clutch – type” lever before measuring. The excitation of the corresponding elections when subjected to IR radiation is reflected in the spectrum as absorption bands at wavelength range from 4000 - 400 cm-1 at scanning speed of 2 cm-1 (Figure 13).
Fig. 13: ATR-Fourier Transform Infrared Spectrophotometer
2.6.3. Scanning Electron Microscope (SEM) and Energy Dispersive Analysis
X-ray (EDX):
Scanning Electron Microscopy (SEM) Model Quanta 250 FEG (Field Emission Gun) attached with EDX Unit (Energy Dispersive X-ray Analyses), with accelerating voltage 30 KV, a magnification of 14x up to 1000000, Gun.1n. FEI Company, Netherlands (Figure 14) is used to examine the surface morphology of the different prepared composite powders. The investigated samples are coated with gold (conductive layer) before imaging using EMITECH K550k sputter coater England. EDX analyzer is used to detect the chemical composition of the synthesized powders. The EDX system has a super ultra-thin window which means that it can analyze a wide range of elements.
Fig. 14: Scanning Electron Microscope with EDX
2.7. Adsorption studies
An accurately 0.02 g of CACs is added into 10 mL of solutions in a 25.0 mL glass tubes containing the specified concentrations of metal ions. The mixtures are shaken at 25.0 oC for a fixed period (2 hrs.) and at the end of shaking periods, the contents are filtered through filter paper. The filtrate is analyzed for final metal concentration using Perkin Elmer 2380 Atomic Absorption Spectro Photometer. Each experiment is performed in triplicate and the average of the results is recorded.
The effects of contact time (5-120 min.) and metal ion concentration on the sorption process are realized using the same methodology. The amounts of metal ion adsorbed onto CACs sorbent, qe (mg g–1), are calculated using the Equation:
qe = (Co – Ce) V/ m
where Co and Ce (mg L–1) are initial and equilibrium concentrations of metal ions, m (g) is the weight of sorbent in the solution and V (L) is the volume of the solution.
The efficiency of adsorption (removal %) is calculated according to the Equation:
% Removal = [(Co – Ce) /Co] × 100
2.7.1. Effect of contact time:
To study the effect of contact time on the adsorption efficiency of Pb2+ and Fe3+ ions by CACs, 0.02 g of CACs is added into 10 mL of 100 mg L-1 Pb2+ (pH 5.7 and 4) and Fe3+ ( pH 2.6) solutions at 25±1°C with time interval from 5 -120 minutes.
2.7.2. Effect of pH:
The effect of pH on the adsorption is performed only for Pb2+. The study is achieved with two pHs values (4 and 5.7) the original solution of lead is at pH 5.7 and to get pH 4 we use HNO3 solution (0.01 M). This process is conducted at 30 minutes contact time, 20 mg dosage of the different composites and 100 ppm of metal solution at 25±1°C. With respect to Fe3+ solution, the experimental work is carried out at pH =2.7 only. This is due to the precipitation of Fe3+ at pH ≥ 3.
2.7.3. Effect of dosage:
To study the effect of dosage (10, 15, 20 mg) on the adsorption efficiency of Pb2+ on CA-Np (as a standard model). The previous dosages are added to 10 mL of 100 mg L-1 Pb2+ at 25±1°C with 30 minutes contact time.
2.8. Adsorption kinetics:
Adsorption kinetic studies are important since they describe the solute uptake rate which controls the residence time of adsorbate at the solid–liquid interface and also provide valuable insights into the reaction pathways.
Pseudo–first–order (Equation I) and pseudo–second–order (Equation II) models were applied in order to investigate the adsorption kinetics of Pb2+ and Fe3+ ions onto CACs. The conformity between experimental data and the model-predicted values is expressed by the correlation coefficients (R2). Meanwhile, the capacity values calculated from the pseudo–first and second–order models are compared with that obtained from the experimental data. The kinetic models can be presented as follows,
ln (qe –qt ) = ln qe – k1 t (I)
t/qt = 1/k2qe2 + t/qe (II)
where qt is the amount of metal ion adsorbed (mg g-1) at time (t), qe is the maximum adsorption capacity (at equilibrium) (mg g-1) and k1(min−1) is the rate constant for pseudo–first–order sorption, qe is the maximum adsorption capacity (at equilibrium) (mg g-1) and k2 (g mg-1 min-1) is the rate constant for the second-order sorption. from plots of t/qt versus t for the second–order reactions, the k2 and qe values are calculated by using the values of intercept and slope.
2.9. Adsorption isotherm models:
The adsorption equilibrium is investigated by using two well-known isotherm models (Langmuir and Freundlich) to provide the fundamental physicochemical data and investigate the applicability of sorption process at a fixed temperature. The equilibrium conditions of the adsorption process are described by utilizing the linearized equations indicated below,
Ce/qe =Ce/qm + 1/ KL qm
where qm (mg g–1) and KL (L mg–1) are constants in Langmuir’s equation which are referred to the maximum adsorption capacity and the Langmuir model constant that is indirectly related to the energy of adsorption. Also qe and Ce parameters represent the equilibrium adsorption capacity and the equilibrium concentration heavy metal ion, respectively.
Freundlich adsorption isotherms assumes aheterogeneous surface with anon-uniform distribution of heat of adsorption over the surface with the possibility of the multilayer adsorption [178, 179].the Freundlich equation which is expressed as
qe =log KFCe1/n
where KF is the measure of adsorption capacity and n is the adsorption intensity linear form of Freundlich
log qe =log KF + (1/n) log Ce
where qe is the amount adsorbed (mg/g), Ce is the equilibrium concentration of adsorbate (mg/L), and KF and n are the Freundlich constants related to the adsorption capacity and adsorption intensity, respectively. A plot of log qe vs. log gives a linear trace with a slope of 1/n and intercept of log KF.
3. Results and Discussion
3.1. characterization of the synthesized composites.
3.1.1. Fourier Transform Infrared (FTIR) analysis:
The Fourier transform infrared (FTIR) spectra of calcium alginate and calcium alginate composites are given in Figures (15-18). Also their frequencies of the absorbtion bands of FTIR spectra and their structure assignments are given in Table (10). Spectrum of calcium alginate (Figure 15), showed important absorption bands regarding hydroxyl, ether and carboxylic functional groups. Stretching vibrations of O–H bonds of alginate appeared in the range of 3000–3600 cm-1 particualy at 3444 cm-1. Stretching vibrations of aliphatic C–H are observed at 2920–2850 cm-1. The observed bands at 1632 and 1454 cm-1 are attributed to asymmetric and symmetric stretching vibrations of carboxylate salt ion, respectively. 1154 cm-1 (CO-stretching of ether group) and 1025 cm-1 (C-O stretching of alcohol group)[180].
The spectra of CAG (Fig.16), the absorption band at around 3442 cm-1 concerned with OH stretching vibration for CA slightly broadened and shifted to a lower wave number with the blending with gelatin, suggesting the formation of an intermolecular hydrogen bond [181]. The strong absorption band at 1631 cm-1 for CA assigned to the asymmetric stretching vibration of COO- has coupled with the absorption band at 1631 cm-1 in gelatin.
CA raw material, reveals asimilar spectra with CAB as shown in Fig 17, in addition some peaks are found at 3699 and 3444 cm-1 are due to lattice OH and bound water stretching vibrations. A strong and sharp band is detected at 1022 cm-1 which is related to Si–OH stretching vibrations [182]. Peaks found at 1384 cm-1 is due to CO3 stretching of calcite, 1034 cm-1 assigned to Si-O stretching, and 875 cm-1 is due to OH bending of the Al-Al-OH group. A similar OH bending vibration is observed for Al-Mg-OH at 842 cm-1, 690 cm-1 assigned to quartz. Also, there is a shoulder peak at 520 cm-1 (Al-O-Si bending), and 464 cm-1 (Si-O-Si bending)[183].
The IR spectra of CAMK and CAD composites revealed a similar spectra with CAB composite spectra (Figure 17). CAD showed a strong band at 1084 and 1048 cm-1 due to Si–OH and Si-O vibrations. These bands overlabed with the band at 1025 cm-1 due to CO-stretching of alcohol group in alginate. A strong and sharp peak at 470 cm-1 are also detected due to Si-O-Si bending. The band of Si-O-Si bending in CAD is the strongest and sharpest compared to CAMK and CAB.
The HAP-Alginate samples (CAHA(I) and CAHA(II)) (Figure 18) revealed a similar spectra, at 1625-1630 cm-1 and 3440-3445 cm-1 (due to the presence of free water), 1454 and 874 cm-1 ( due to CO32- ions), 3570 and 630-633 cm-1 (due to structural OH of hydroxy apatite). These peaks due to the hydroxyapatite phase [184, 185]. The most intensive bands in the range of 1044 –1090 cm-1 corresponded to the triply degenerated asymmetric stretching vibrations of P-O. Otherwise the peak at 962.97 cm-1 indicates the non-degenerated asymmetric mode of PO43-. The very strong and sharp bands observed at 569-572 and 602 -603 cm-1 attributed to triply degenerated bending mode of the O-P-O in PO43- group. The larges parting distance of these bands revels the crystalline phase [184, 185].
Table (10): Assignments of the absorption bands of the IR spectra λ (cm-1) of the prepared composites
Peak Assignment Strength CA CAB CAMK CAHA(I) CAHA(II) CAD CAG
Structural OH of addittives w - 3699 - 3571 3570 - -
OH stretching mode of adsorbed water molec. or OH of alginate lattice structure s , b 3444 3444 3444 3444 3444 3443 3442
Stretching mode of aliphatic C–H. w 2924 2923 2924 2923 2924 2924 2924
asymmetric stretching vibrations of carboxylate salt ion m 1631 1634 1631 1629 1630 1629 1631
symmetric stretching vibrations of carboxylate salt ion vw 1427 1461 1431 1433 1419 1433 1427
CO-stretching of ether group vw 1155 1150 1153 - - 1103 1155
CO-stretching of alcohol group w 1024 1033 - - - - 1024
Si–OH stretching vibrations mode w - 1080 1081 - - 1084
CO3 stretching of calcite w 1384 1384 1384 1384 1384 1384 1384
Si-O stretching w - 1033 1052 - - 1048 -
OH bending of the Al-Al-OH group vw - 875 850 - - 845 -
OH bending vibration of Al-Mg-OH group of quartz vw - 842
690 777 - - 797 -
Al-O-Si bending vw - 520 510 - - 526 -
bending vibration mode of Si-O-Si s - 464 456 - - 470 -
vibration mode CO32- vw - - - 1454
874 1457
874 - -
structural OH in HA w - - - 631 631 - -
triply degenerated asymmetric stretching vibrations of P-O vs - - - 1044
1090 1048
1089 - -
non-degenerated asymmetric of PO43- vw - - - 962 962 - -
triply degenerated bending mode of PO43- ms - - - 602
569 602
571 - -
s = strong w = weak vw = very weak b = broad m = meadium
Figure 15: FTIR Spectra of Calcium alginate.
Figure 16: FTIR Spectra a- CA, b- CAG
Figure 17: FTIR Spectra a- CA, b- CAB, c- CAMK ,d- CAD.
Figure 18: FTIR Spectra a- CAHA(I) , b-CAHA(II)
3.1.2. X-ray diffraction for crystal phase detection:
Figure(19) presents the X-ray diffraction pattern of CA and CA composites in the range of 2θ = 5-60o diffraction degree . Two typical peaks in 2θ =16° and 22° are observed for calcium alginate. The XRD of CAG shows typical peaks around 12° and 21°[186].
The samples of CAB, CAMK and CAD powders revealed very similar XRD pattern peaks of quartz phase. These composites showed the maximum relative intensity (I/I0) peak of 100% quartz in the rang of 2θ =26.65o and 20.85o and d (Ao) value spacing equal 3.34 and 2.8, (reference code :01-070-3755). On the other hand, additional peaks are detected due to the presence of minor amounts of calcite (CaCO3) and halite (NaCl). The presence of calcite phase may be attributed to carbonation of calcium during composite synthesis, while the presence of NaCl may be due to the reaction of Na+ ions of alginate with Cl- ions of CaCl2 solution and its trapping bettween the layers, which can not completely removed by washing.
The well resolved XRD peaks of The sample HAP(I) ,CAHA(I), HAP(II) and CAHA(II) could be easily indexed on the basis of hexagonal crystal system of space group P63/m with respect to JCPDS file no. 9-432. They also revealed very similar XRD pattern peaks of 100 % HAP (reference code :01-086-1194) in the rang of 2θ =31o and 32o and d (Ao) value spacing equal 2.81 and 2.78. There is no any considerable shifts in 2θ are detected between HAP and HAP composite . The diffraction peaks of HA and CA-HAP exhibit sharp diffraction peaks which indicate the high crystallinity of the structure and there is no any additional phases are detected.
XRD analysis of HAP(II) and CAHA(II) also indicated the absence of secondary phases, such as tri calcium phosphate (TCP) or calcium oxide (CaO). In the case of HAP(I) and/or CAHA(I), their diffraction patterns revealed additional phase of β-tricalcium phosphate, beside the HAP as a main phase.
Hydoxyapatite prepared from bovine bone contain certain amount of carbonate (CO32-) in its lattice structure [175]. So the sample CAHA(I) are expected to be a carbonated apatite type, and carbonate ions affect on the the degree of crystallinite. For this reason, CAHA(II) ( HA prepared from bovine bones) revealed higher crystal size (83 nm) than CAHA(I) which prepared from eggshells (59 nm). There is no significant differences in the crystal sizes of the prepared HAP and HAP Compsites. Table (11) showed the crystal size, 2θ, d-spacing(oA), maximum relative intensity (I/Io) peak and the main phase detected for the prepared adsorbents
Table (11): 2θ, d-spacing(oA), Crystal size ( nm ) of the maximum relative intensity (I/Io) peak and the main phase detected for the prepared adsorbents.
Character
2θ
d-spacing (Ao)
Crystal size (nm)
Main phase detected
CAHA(I)
31.83
2.81
59
HA
CAHA(II)
31.77
2.81
83
HA
CAB
26.65
3.34
83
Quartz
CAMK
26.68
3.34
84
Quartz
CAD
26.63
3.34
59
Quartz
Figure 19: X-ray diffraction pattern of the prepared calcium alginate composite powders.
Figure 19: continue
Figure 19: continue
Figure 19: continue
3.1.3. Scanning Electron Microscope (SEM) and Energy Dispersive X-ray (EDX)
The SEM images and EDX elemental analysis of synthesized calcium alginate and calcium alginate composites are shown in Figure (20). The images are taken at 300× and 2500× magnification. SEM investigations showed the presence of agglomerates of irregular shape particles.
According to the results obtained by EDX analysis, carbon, oxygen and calcium are the major constituents of all the prepared composites related to calcium alginate polymer present, Also, Almost prepared composites confirmed the presence of calcite and halite as additional phases in minor amounts which is in aggreement with X-ray results.
EDX analyses of CA and CAG revealed that the major elements present are carbon and oxygen corresponding to polymer composition. The SEM of CAG (Fig.20(i)) showed a smooth and homogeneous morphology, suggesting high miscibility and blend homogeneity between calcium alginate and gelatin.
In the case of CAB, CAMK and CAD, the presence of an obvious peak related to the Si compounds is evident. According to the results obtained by EDX analysis, weigh percent of elements present indicating a large portion of the composites is composed of Si compounds which is suitable for an efficient sequestering metal cations from aqueous solution[129].
EDX analyses of the prepared HAP revealed that inorganic phases of bovin bone and egg shells were mainly composed of calcium and phosphorus as the major constituents with some minor components such as C, O, Na, Mg and Si. The weight and atomic percentage shows that the Ca/P ratio around 1.7 and 1.8 which is below 2 and acceptable where the ideal Ca/P ratio of HA is 1.67[174], in HAP composites (CAHA(I) and CAHA(II)) this ratio increased than 2 due to the excess amount of Ca crosslinkage of calcim alginate presents.
a) CA
Element Wt % At %
C K 35.08 46.27
O K 45.62 45.17
Na K 2.08 1.43
Cl K 6.18 2.76
Ca K 11.05 4.37
Total 100 100
b) CAB
Element Wt % At %
C 16.29 25.90
O 37.59 44.87
Na 1.14 0.95
Al 13.61 9.77
Si 16.26 11.06
Cl 6.83 3.68
Ca 7.08 3.37
Ti 1 0.4
total 100 100
Figure 20: SEM images and EDX analysis of prepared samples a) CA & b) CAB .
c) CAMK
Element Wt % At %
C K 9.75 16.10
O K 41.60 51.55
Na K 1.07 0.93
Al K 4.43 3.26
Si K 26.61 20.91
Cl K 6.25 3.50
Ca K 5.85 2.90
K K 0.98 0.49
Mg 0.45 0.37
Total 100 100
d) CAD
Element Wt % At %
CK 27 39.27
OK 39.73 43.39
Na K 2.46 1.87
Al K 2.30 1.49
Si K 5.16 3.21
Cl K 7.79 3.84
Ca K 14.62 0.21
K K 0.46 0.21
Mg 0.47 0.34
Total 100 100
Figure 20: SEM images and EDX analysis of prepared samples c) CAMK & d) CAD.
e) HAP(I)
Element Wt % At %
C K 4.66 9.69
O K 29.05 45.31
P K 20.45 16.47
Ca K 45.84 28.53
Total 100 100
f) HAP(II)
Element Wt % At %
C K 3.96 8.22
O K 28.88 45.05
Na K 2.42 2.62
Mg K 0.76 0.78
P K 18.95 15.27
Ca K 45.03 28.04
Total 100 100
Figure 20: SEM images and EDX analysis of prepared samples e) HAP(1) & f) HAP(2).
g) CAHA(I)
Element Wt % At %
C K 20.69 33.23
O K 36.42 43.90
Na K 1.74 1.46
P K 10.18 6.34
Cl K 2.62 1.42
Ca K 28.35 13.64
Total 100 100
h) CAHA(II)
Element Wt % At %
C K 7.32 14.25
O K 31.95 46.71
Na K 0.70 0.71
Mg k 0.40 0.39
P K 17.06 12.86
Cl K 2.71 1.79
Ca K 39.86 23.27
Total 100 100
Figure 20: SEM images and EDX analysis of prepared samples. g) CAHA(I) & h) CAHA(II).
i) CAG
Element Wt % At %
C K 39.21 49.41
O K 44342 44.07
Na K 2.08 1.43
Cl K 3.15 1.8
Ca K 11.05 4.37
Total 100 100
Figure 20: SEM images and EDX analysis of prepared samples. i) CAG
3.2. characterization of the composites after metal ion uptake:
3.2.1. FTIR Analysis after metal ion uptake:
FTIR spectra of CA and CAHA(I) after targeted metal ions uptake ( Pb2+, Fe3+ ) are shown in Figures 21 and 22 respectively and their peak assignments are represented in Table (12). The result of FTIR spectra showed that there is no new absorption peaks were detected . There are alittle peak shifts which may be attributed to the corporation and substitution of metal ion in the lattice structure of CA and CAHA(I).
Table 12: Change in the absorption bands of the IR spectra λ (cm-1) of CA and CAHA(I) powder after metal ion uptake
Peak Assignment Strength CA
CA
+
Pb2+
CA
+
Fe3+ CAHA(I)
CAHA(I)
+
Pb2+
CAHA(I)
+
Fe3+
Lattice Structural OH of hydroxyapatite w -
- - 3571 3571 3571
OH stretching mode of adsorbed water molec. or OH of alginate lattice structure s , b 3444 3444 3434 3444 3444 3444
Stretching mode of aliphatic C–H . w 2924 2924 2923 2923 2924 2923
asymmetric stretching vibrations of m 1631 1598 1631 1629 1599 1615
symmetric stretching of COO- ion vw 1427 1425 1424 1433 1433 1433
CO-stretching of ether group. vw 1155 1156 1155 -
CO-stretching of alcohol group w 1024 1021 1021 -
CO3 stretching of calcite w 1384 1384 1383 1384 1384 1381
vibration mode CO32- vw - 1454
874 1454
874 1458
874
structural OH in HA w - 631 631 631
triply degenerated asymmetric stretching vibrations of P-O vs - 1044
1090 1045
1090 1045
1090
non-degenerated asymmetric of PO43- vw - 962 962 962
triply degenerated bending mode of PO43- ms - 602
569 602
570 601
568
s = strong w = weak vw = very weak b = broad m = meadium
Figure 21: FTIR Spectra a- CA , b-CA+ Pb2+ , c- CA + Fe3+
Figure 22: FTIR Spectra of a- CAHA(I) , b- CAHA(I)+ Pb2+ , c- CAHA(I)+ Fe3+
3.2.2. X-ray diffraction for crystal phase detection after metal ion uptake:
XRD patterns of CA and CAHA(I) after Pb2+ and Fe3+ metal ions removal did not revealed any new phases. supported the proposal that Pb2+ and Fe3+ ions uptake was not dependent on dissolution/precipitation mechanisms. Pb2+ and Fe3+ ions removal may be occurs by adsorption mechanisms like surface complexation or ionic exchange [187]. As it can be seen in Figures 23-24, XRD patterns showed some changes in their relative intensities and crystal sizes (Table 13). Also, ther are some little shifts in d- spacing values, this may be due to the ion exchange between Ca2+ and metal cations of Pb2+ and Fe3+ in lattice structure of CA and CAHA(I) nanopowders.
Table 13: 2θ, d-spacing(Ao), Crystal size ( nm ) of the maximum relative intensity (I/Io) peak of CAHA(I) before and after metal ion uptake.
Character
2θ d-spacing (Ao) Crystal size (nm)
CAHA(I)
31.83 2.811 59
CAHA(I) + Pb2+
31.74 2.819
60
CAHA(I) + Fe3+ 31.71 2.821 37
Figure 23: X-ray diffraction pattern of CA after metal ion uptake
Figure 24: X-ray diffraction pattern of CAHA(I) after metal ion uptake
3.2.3. Scanning Electron Microscope (SEM) and Energy Dispersive X-ray (EDX) after metal ion uptake:
SEM of CA and CAHA(I) after metal ion uptake revealed some changes in morphology and microstructure(Fig. 26,27). Morever EDX results indicated the presence of Pb2+ and Fe3+ ions with CA and CAHA(I). As can seen in Figures 26 and 27, the peak of Fe3+ is more intense than that of Pb2+ and this confirms, the higher value of Fe3+ ions uptake by CA and CAHA(I) comparing with Pb2+ ions. The weight and atomic percentage of Ca2+ ions in CA and CAHA(I) after metal ion uptake was less than its value before metal ion removal as listed in Table (14). The decrease in calcium percentage may be attributed to ion exchange between targeted metal ions and; (1)calcium ions of calcium alginate in CA and CAHA(I) as follows:
Ca(ALG)2 + Pb2+ Pb(ALG)2 + Ca2+
(2) or Ca ions of HA present in CAHA(I). This ion exchange mechanism between Pb2+ ions (as example) and Ca2+ ions of HA produced anew phase of hydroxypyromorphite [132, 188], this mechanism is expressed as:
Ca10 (PO4)6 (OH)2 + x Pb2+ x Ca2+ + Ca10-xPbx (PO4)6(OH)2
However , this phase is not detected in the present study this may be attributed to under limit of XRD
Table 14: Ca2+ ion percentage in CA and CAHA(I) before and after metal ion uptake.
Character
Ca2+ % before uptake
Ca2+ % after uptake
Wt %
+Pb2+ (Wt %)
+ Fe3+ (Wt %)
CA 11.05 5.69 1
CAHA(I) 28.35 15.61 23.43
Also the uptake process may be occurs during chelation bonding of targeted metal ions with two carboxylic groups of alginate and one or two OH sites of the alginate ring (Fig 25) [189]. In this case metal ions may forms complexes with two adjacent alginate rings. Here,‘‘adjacent’’ means either two neighbor alginate rings of a single polymeric chain (intramolecular chelation) or two rings from two parallel chains (intermolecular chelation) [189].
Figure 25: potential active sites of CA which may bonds with targeted metal ion
a) CA + Fe+3
Element Wt % At %
C K 15.77 32.25
O K 27.57 42.32
Si K 0.41 0.35
Cl K 0.62 0.43
Ca K 1 0.61
Fe K 54.64 24.03
Total 100 100
b) CAHA(I) + Fe+3
Element Wt
% At
%
C K 7.60 12.77
O K 36.54 55.02
Al k 0.37 0.33
P K 14.03 12.91
Ca K 23.43 11.52
Fe K 18.03 7.78
Total 100 100
Figure 26: SEM images and EDX analysis of CA and CAHA(I) after iron removal
a) CA + Pb+2
Element Wt % At
%
C K 32.44 51.66
O K 36.01 43.05
Al k 0.30 0.21
Pb M 25.56 2.36
Ca K 5.69 2.71
Total 100 100
b) CAHA(I) + Pb+2
Element Wt % At %
C K 9.53 18.79
O K 32.94 48.74
Al k 0.37 0.32
P K 15.46 11.81
Pb M 8.99 1.03
Ca K 32.71 19.31
Total 100 100
Figure 27: SEM images and EDX analysis of CA and CAHA(I) after lead removal.
3.3. Metal ion uptake by adsorption process.
3.3.1. Effect of contact time on the adsorption process.
3.3.1.1 Calcium Alginate (CA):
The effect of contact time on the adsorption capacity of CA for Pb2+ ( natural pH of 5.7), Pb2+ (pH 4) and Fe3+ ( natural pH of 2.6) is represented in Figure 28. The equilibrium points of adsorption were attained within the first 60 min. for Pb2+ (82.3%) , 30 min. for Pb2+ (86.82%) and 30 min. for Fe3+ (94.15%) of contact time. the adsorption of Pb2+, Pb2+( pH 4) and Fe3+onto CA was rapid in the initial stages up to 30 min, and was almost same at high contact time. This behavior could be attributed to the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption process. The exponential phase for Pb2+( pH 4) and Fe3+ was found to be between 5 minutes and 30 minutes. For Pb2+ it was found to be between 5 minutes and 60 minutes. Further increase in contact time led to no significant adsorption of metal ions by CA probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches its maximum after 120 minutes (94.18%).
Figure 28: Effect of contact time on the adsorption of Pb2+( pH 5.7) , Pb+2(pH 4) and Fe3+ ions(pH 2.6) from aqueous solution by CA under experimental conditions of CA mass 0.02 g/10mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.1.2 CAB
The effect of contact time on the adsorption capacity of CAB for Pb2+ ( natural pH of 5.7), Pb2+(pH 4) and Fe3+ ( natural pH of 2.6) is represented in Figure 29. The equilibrium points of adsorption were attained within the first 60 minutes for Pb2+ (66.40%), 30 minutes for Pb2+ (pH 4)(68.2%) and 30 minutes for Fe3+ (89.42%) of contact time. the adsorption of Pb2+, Pb2+( pH 4) and Fe3+onto CAB was rapid in the initial stages up to 30 min, and was almost same at high contact time. This behavior may be due to the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption process. The exponential phase for Pb2+( pH 4) and Fe3+ was found to be between 5 minutes and 30 minutes. For Pb2+ it was found to be between 5 minutes and 60 minutes. The sorption efficiency by CAB was almost constant probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches its maximum after 120 minutes (89.6%).
Figure 29: Effect of contact time on the adsorption of Pb2+(pH=5.7)Pb+2(pH=4) and Fe3+(pH= 2.6)ions from aqueous solution by CAB under experimental conditions of CAB mass 0.02 g/10 mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.1.3 CAMK
The effect of contact time on the removcal of Pb2+ ( natural pH of 5.7), Pb2+ ( pH 4) and Fe3+ ( natural pH of 2.6) by CAMK is represented in Figure 30. The adsorption percent of metal ions was fast at initial stages and gradually become slower until the equilibrium is attained. The optimal contact time to attain equilibrium was experimentaly found to be about 30 min. for Pb2+ (69.26%) , Pb2+( pH 4) (77.7%) and Fe3+ (89.87%). the adsorption of Pb2+, Pb2+( pH 4) and Fe3+onto CAMK was rapid in the initial stages up to 30 minutes, and was almost same at high contact time. This behavior could be because of the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption processThe exponential phase for Pb2+, Pb2+( pH 4) and Fe3+ was found to be between 5 minutes and 30 minutes. Further increase in contact time led to no significant adsorption of metal ions by CAMK probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches its maximum after 120 minutes (90%).
Figure 30: Effect of contact time on the adsorption of Pb2+ (pH= 5.7), Pb+2( pH= 4) and Fe3+ (pH=2.6) ions from aqueous solution by CAMK under experimental conditions of CAMK mass 0.02 g/10 mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.1.4. CAHA(I)
The effect of contact time on the adsorption capacity of CAHA(I) for Pb2+ (natural pH of 5.7), Pb2+ (pH 4) and Fe3+ ( natural pH of 2.6) is represented in Figure (31). As it can be seen in fig.22, the equilibrium points of adsorption were attained within the first 60 minutes for Pb2+ (80.78%), 30 minutes for Pb2+ (pH 4) (86.62%) and 30 minutes for Fe3+ (99.33%) of contact time. the adsorption of Pb2+, Pb2+( pH 4) and Fe3+onto CAHA(I) nanopowders was rapid in the initial stages up to 30 min, and was almost same at high contact time. This behavior could be attributed to the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption process. The exponential phase for Pb2+( pH 4) and Fe3+ was found to be between 5 minutes and 30 minutes. For Pb2+ it was found to be between 5 minutes and 60 minutes. The sorption efficiency by CAHA(I) was almost constant probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches its maximum after 120 minutes (99.34%)
Figure 31: Effect of contact time on the adsorption of Pb2+ (pH= 5.7), Pb+2(pH= 4) and Fe3+ ions (pH= 2.6) from aqueous solution by CAHA(I) under experimental conditions of CAHA(I) mass 0.02 g/10 mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.1.5. CAHA (II)
The adsorption efficiency of Pb2+ ( natural pH of 5.7), Pb2+ (pH 4) and Fe3+ ( natural pH of 2.6) by using CAHA (II) was tested at different contact time. Asit can be seen in figure 32, the equilibrium points of adsorption were attained within the first 30 minutes for Pb2+ (74.08%) and Pb2+ (pH 4) (84.67%) while for Fe3+ were attained within the first 10 minutes (98.65%) of contact time. the adsorption of Pb2+, Pb2+( pH 4) and Fe3+onto CAHA(II) nanopowder was rapid in the initial stages up to 30 min, and was almost same at high contact time. This behavior could be attributed to the availability of a large number of active sites for rapid surface metal ions binding during the initial stages of adsorption process. The exponential phase for Pb2+and Pb2+( pH 4) was found to be between 5 minutes and 30 minutes. For Fe3+ it was found to be between 5 minutes and 10 minutes. Further increase in contact time led to no significant adsorption of metal ions by CAHA (II) probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches 98.78% after 60 minutes and its maximum after 120 minutes (98.81%).
Figure 32: Effect of contact time on the adsorption of Pb2+ (pH= 5.7), Pb+2(pH= 4) and Fe3+ ions (pH= 2.6) from aqueous solution by CAHA (II) under experimental conditions of CAHA (II) mass 0.02 g/10 mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.1.6. CAD
The effect of contact time on the adsorption capacity of CAD for Pb2+ ( natural pH of 5.7), Pb2+(pH 4) and Fe3+ ( natural pH of 2.6) is represented in Figure 33. The equilibrium points of adsorption were attained within the first 30 minutes for Pb2+ (68.96%), Pb2+( pH 4) (63.23%) and Fe3+ (86.72%) of contact time. The adsorption of Pb2+, Pb2+(pH 4) and Fe3+onto CAD was rapid in the initial stages up to 30 min, and was almost same at high contact time. This behavior could be attributed to the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption process. The exponential phase for Pb2+, Pb2+( pH 4) and Fe3+ was found to be between 5 minutes and 30 minutes. The sorption efficiency by CAD was almost constant probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches its maximum after 120 minutes (86.95%).
Figure 33: Effect of contact time on the adsorption of Pb2+ (pH= 5.7), Pb+2(pH= 4) and Fe3+ ions (pH= 2.6) from aqueous solution by CAD under experimental conditions of CAD mass 0.02 g/10 mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.1.7. CAG
The sorption efficiency exhibited by of CAG for Pb2+ ( natural pH of 5.7), Pb2+ (pH 4) and Fe3+ ( natural pH of 2.6) is depicted in Figure 34. The equilibrium points of adsorption were attained within the first 30 minutes for Pb2+ (83.43%), Pb2+( pH 4) (87.93%) and Fe3+ (91.11%) of contact time. The adsorption of Pb2+, Pb2+( pH 4) and Fe3+onto CAG was rapid in the initial stages up to 30 min, and was almost same at high contact time. This behavior could be due to the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption processThe exponential phase for Pb2+, Pb2+( pH 4) and Fe3+ was found to be between 5 minutes and 30 minutes.. The sorption efficiency by CAG was almost constant probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches 91.55% after 60 min. and reaches its maximum after 120 minutes (91.81%).
Figure 34: Effect of contact time on the adsorption of Pb2+ (pH= 5.7), Pb+2(pH 4) and Fe3+ ions (pH= 2.6) from aqueous solution by CAG under experimental conditions of CAG mass 0.02 g/10 mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.2. Effect of pH
The effect of pH on the adsorption is performed only for Pb2+ because of Fe3+ solution was stable only at pH lower than 3. The study is achieved with two pHs values (4 and 5.7) the original solution of lead was at pH=5.7 and pH=4 at 5-120 minutes contact time, 20 mg dosage of the different composites and 100 ppm of metal solution at 25±1°C. As can be seen in Figures 28-34, the adsorption efficiency of Pb2+ at pH=4 is higher than that of pH=5.7 for all the composites.
3.3.3. Effect of adsorbent dosage on metal ion adsorption.
The experimental results of the adsorption of Pb2+ on CA ( as astandard model ) as afunction of adsorbent dosage 10, 15 and 20 mg/10 mL, initial Pb2+ concentration of 100 mgL-1, natural pH of 5.7, temperature 25oC at the optimal contact time (30 min) and interval contact time ( 5-30 min) are shown in fig. 36 and 37 respectively. As can be seen in Figures 35 and 36, the Pb2+ adsorption percent rapidly increased with the increase in the adsorbent dosage . this can be attributed to higher adsorbent dosage due to the increased surface area providing more adsorption sites available which gave rise to higher removal of lead [190].
Figure 35: Effect of CA dosage on Pb2+ at contact time 30 min. , initial concentration of 100 mgL-1, natural pH of 5.7 and temperature 25±1oC.
Figure 36: Effect of CA dosage on Pb2+ at contact time 5-30 min. , initial concentration of 100 mgL-1, natural pH of 5.7 and temperature 25±1oC.
3.4. kinetics studies of the adsorption process.
The kinetic study is useful to predict the adsorption rate which is very important in modeling and designing of the adsorption process [191]. The pseudo-first rate equation of lagergren and pseudo-second order kinetic model, as the most widely used models, are used to evaluate the mechanism of adsorption process.
3.4.1. The pseudo first-order model:
The linear form of the pseudo first-order kinetic rate equation is given as follows:
ln (qe –qt ) = ln qe – k1 t
where qt is the amount of metal ion adsorbed (mg g-1) at time (t), qe is the maximum adsorption capacity (at equilibrium) (mg g-1) and k1(min−1) is the rate constant for pseudo–first–order sorption, qe is the maximum adsorption capacity (at equilibrium) (mg g-1). The kinetic of adsorption are evaluated at an initial concentration of 100 mg/L for Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+(pH=2.6), adsorbent dosage of 0.02 g/10 mL and temperature of 25oC. Plot of ln (qe –qt ) vs t is drawn as shown in Figures 37-43.
The rate constant at equilibrium (qe) and regression coefficient (R2) obtained from the plots of pseudo-first rate equation of adsorbed Pb2+ and Fe3+ at equilibrium (qe) for all the adsorbents are given in Tables 15,16 and 17 respectively. As it can be seen in Figures (37-43) and Tables (15-17), The regression coefficient does not close to unity. Also, the values of qe obtained from pseudo-first order equation for all the adsorbent are different and not matched notably with the experimental qe value.
Figure 37: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CA at experimental conditions of 100 mg/L metal ion concentration, CA mass 0.02 g/10 mL and 25±1oC.
Figure 38: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAB at experimental conditions of 100 mg/L metal ion concentration, CAB mass 0.02 g/10 mL and 25±1oC.
Figure 39: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAMK at experimental conditions of 100 mg/L metal ion concentration, CAAMK mass 0.02 g/10 mL and 25±1oC.
Figure 40: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+( pH= 4) and Fe3+ (pH=2.6) adsorbed on CAHA(I) at experimental conditions of 100 mg/L metal ion concentration, CAHA(I) mass 0.02 g/10 mL and 25±1oC.
Figure 41: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAHA(II) at experimental conditions of 100 mg/L metal ion concentration, CAHA(II) mass 0.02 g/10 mL and 25±1oC.
Figure 42: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+(pH=2.6) adsorbed on CAD at experimental conditions of 100 mg/L metal ion concentration, CAD mass 0.02 g/10 mL and 25±1oC.
Figure 43: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAG at experimental conditions of 100 mg/L metal ion concentration, CAG mass 0.02 g/10 mL and 25±1oC.
3.4.2. The pseudo second-order model:
The linear form of the pseudo second-order kinetic rate equation is given as follows:
t/qt = 1/k2qe2 + t/qe
where qt is the amount of metal ion adsorbed (mg g-1) at time (t), qe is the maximum adsorption capacity (at equilibrium) (mg g-1) and k2 (g mg-1 min-1) is the rate constant for the second-order sorption. from plots of t/qt versus t for the second–order reactions Figures 44-50, the k2 and qe values are calculated by using the values of intercept and slope as summarized in Tables (15, 16, 17).
The rate constant and regression coefficient (R2) obtained from the plots of pseudo-second rate equation of adsorbed Pb2+ (pH=5.7), Pb2+( pH 4) and Fe3+ (pH=2.6) at equilibrium (qe) for all the adsorbents are given in Tables 15,16 and 17 respectively . from the linear plots, the qe,experimental and the qe,calculated values are very close to each other, and also, the calculated coefficients of determination, R2, are close to unity.
Table 15: Experimental and calculated parameters of pseudo-first and second order kinetic models of Pb2+ (natural pH of 5.7) adsorbed on CA and CACs powder.
Qe experimental Pseudo first order Pseudo second order
qe calculated
K1
R2 qe calculated
K2
R2
CA
41.15
3.412
0.0769
0.9183
41.15
1.91x10-6
0.9997
CAB
33.2
5.77
0.03538
0.4826
33.69
3.4x10-5
0.9923
CAMK 34.89 11.881 0.12773 0.9636 36.16 2.41x10-5 0.9987
CAHA(I) 40.39 16.29 0.11735 0.9438 42.19 2.06x10-5 0.9984
CAHA(II)
37.09 35.55 0.19646 0.8367 39.65 3.21x10-5 0.9951
CAD 34.87 27.91 0.1423 0.8184 38.13 6.22x10-5 0.9778
CAG 42.07 11.189 0.10384 0.6573 43.3 1.33x10-5 0.9981
Table 16 . Experimental and calculated parameters of pseudo-first and second order kinetic models of Pb2+ (pH=4) adsorbed on CA and CACs powder.
qe Experimental Pseudo first order Pseudo second order
qe calculated
K1
R2 qe
calculated
K2
R2
CA 43.45 9.698 0.1718 0.9431 44.09 4.7x10-6 0.9997
CAB
34.55
13.06
0.1026
0.9001
35.95
3.63x10-5
0.9967
CAMK
38.92
27.24
0.1758
0.8219
40.7
2.34x10-5
0.9978
CAHA(I)
43.39
32.13
0.1921
0.9459
45.74
1.53x10-5
0.9955
CAHA(II) 42.45 15.99 0.1938 0.9859 44.3 1.47x10-5 0.9974
CAD
32.21
14.73
0.107
0.9513
34.16
5.62x10-5
0.9944
CAG
44.1
21.03
0.1394
0.6528
44.62
1.18x10-5
0.9932
Table 17: Experimental and calculated parameters of pseudo-first and second order kinetic models of Fe3+ (natural pH of 2.6) adsorbed on CA and CACs powder.
qe experimental Pseudo first order Pseudo second order
qe calculated
K1
R2 qe
calculated
K2
R2
CA
47.09
43.46
0.2330
0.8915
49.26
9.57x10-6
0.9995
CAB
44.75
39.40
0.2171
0.9256
47.12
1.37x10-6
0.9963
CAMK
44.95
29.695
0.2163
0.8390
46.25
6.56x10-6
0.9997
CAHA(I)
49.72
1.48
0.1037
0.8751
50
2.2x10-3
0.9999
CAHA(II)
49.39
1.41
0.0747
0.7308
50
2.2x10-3
0.9999
CAD
43.4
21.32
0.1857
0.8184
44.3
9.63x10-6
0.9974
CAG
45.75
5.47
0.0938
0.6573
46.08
3.18x10-6
0.9981
Figure 44: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4)and Fe3+ (pH=2.6) adsorbed on CA at experimental conditions of 100 mg/L metal ion concentration, CA mass 0.02 g/10 mL and 25±1oC.
Figure 45: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAB at experimental conditions of 100 mg/L metal ion concentration, CAB mass 0.02 g/10 mL and 25oC.
Figure 46: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAMK at experimental conditions of 100 mg/L metal ion concentration, CAMK mass 0.02 g/10 mL and 25±1oC.
Figure 47: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAHA(I) at experimental conditions of 100 mg/L metal ion concentration, CAHA(I) mass 0.02 g/10 mL and 25±1oC.
Figure 48: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAHA(II) at experimental conditions of 100 mg/L metal ion concentration, CAHA(II) mass 0.02 g/10 mL and 25±1oC.
Figure 49: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAD at experimental conditions of 100 mg/L metal ion concentration, CAD mass 0.02 g/10 mL and 25±1oC.
Figure 50: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAG at experimental conditions of 100 mg/L metal ion concentration, CAG mass 0.02 g/10 mL and 25±1oC.
from all the obtained results illustrated in Figures 37-50 and Tables 15-17, it is obvious that the regression coefficient (R2) from pseudo-second order rate equation for all the adsorbents is higher than that of the pseudo-first order model. On the basis of the regression coefficient and calculated values of adsorption capacity, the adsorption process is found to obey and exhibited best fit to the pseudo-second-order kinetic model which mean that the rate-limiting step might be chemical adsorption or chemisorption involving valency forces through exchange of electrons between the sorbate and the sorbent, also only one ion of the metal is sorbed onto two sorption sites on the sorbent surface [192, 193].
3.4.3. Prediction of adsorption rate-limiting step
There are essentially three consecutive mass transport steps associated with the adsorption of solute from the solution by an adsorbent. These are (1) film diffusion, (2) intraparticle or pore diffusion, and (3) sorption into interior sites. The third step is very rapid and hence, film and pore transports are the major steps controlling the rate of adsorption [179, 194].
The most commonly used technique for identifying the mechanism involved in the adsorption process is by fitting an intraparticle diffusion plot [195]. The amount of metal ions adsorbed (qt) at time (t), is plotted against the square root of t (t1/2), according to Eq. proposed by Weber and Morris as follows:
Qt = Kid t0.5 + C
where C is constant and kid is the intraparticle diffusion rate constant (mg/g min1/2), qt is the amount adsorbed at a time (mg/g), t is the time (min), and kid (mg/g min1/2) is the rate constant of intraparticle diffusion. Due to the varying extent of adsorption in the initial and final stages of the experiment two straight lines with different slopes are obtained (Figures 51-57).
The two regions in the qt vs. t0.5 plot suggest that the sorption process proceeds by surface sorption and intraparticle diffusion. The initial rapid uptake can be attributed to the boundary layer effects (film diffusion). After the external surface loading was completed, the intraparticle diffusion or pore diffusion takes place, The second linear part of the plot presented in Fig.51-57, corresponds to the transportation of Pb2+ and Fe3+ within CA and CACs particles [196]. The slope of the second linear portion of the plot has been defined to yield the intraparticle diffusion parameter of ki1, Ri12 (first stage) and ki2, Ri22 (second stage) are listed in Tables (18-20). On the other hand, the intercept of the plot give an idea about the thikness of boundary layer effect [197]. The larger the intercept, the greater the contribution of the surface sorption in the rate-controlling step [195].
As it can be seen in Figures (51-57), the plot indicated that the intraparticle diffusion was not the rate-controlling step because it did not pass through the origin [196]. The deviation of the straight lines from the origin may be due to the difference in the rate of mass transfer in the initial and final stages of adsorption [198]. Further, the first straight portion is attributed to a macropores diffusion process and the second linear portion can be ascribed to a micropore diffusion process [192, 199]. In addition, it is clear from Fig. (51-57) that the first stage is faster than the second one. This behaviour may be correlated with the very slow diffusion of the adsorbate from the surface film into the micropores, which are the least accessible sites for adsorption [197].
Figure 51: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+(pH=2.6) on CA.
Figure 52: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) on CAB composite.
Figure 53: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) on CAMK composite.
Figure 54: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+( pH 4) and Fe3+ (pH=2.6) on CAHA(I) composite.
Figure 55: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) on CAHA(II) composite.
Figure 56: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) on CAD composite.
Figure 57: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) on CAG composite.
Table 18: The parameters of intraparticle diffusion model for adsorption of Pb2+ (pH=5.7) onto CA and CACs.
Intrapartical diffusion modle
Ki1
(mg/g m0.5)
Ki2
(mg/g m0.5)
Ri12
Ri22
Intercept i1
Intercept i2
CA
0.617 0.071 0.9541 0.7161 37.55 40.50
CAB
1.460 0.13 0.8094 0.7094 24.3 32.07
CAMK
1.682 0.049
0.9365 0.6601 25.94 34.31
CAHA(I)
3.53 0.159 0.9989 0.6689 23.43 38.85
CAHA(II)
2.087 0.019 0.9917 0.8957 25.62 36.93
CAD
3.77 0.15 0.7123 59.22 15.48 33.45
CAG
1.36 0.12 0.9615 0.8696 33.92 41.05
Table 19: The parameters of intraparticle diffusion model for adsorption of Pb2+ (pH=4) onto CA and CACs.
Intrapartical diffusion modle
Ki1
(mg/g m0.5)
Ki2
(mg/g m0.5)
Ri12
Ri22
Intercept i1
Intercept i2
CA
1.11 0.066 0.9056 0.4133 38.07 42.87
CAB
1.71 0.10 0.9716 0.7251 33.6 24.54
CAMK
1.88 0.021 0.9745 0.9251 28.48 38.74
CAHA(I)
3.21 0.086 0.8603 0.3367 27.94 42.06
CAHA(II)
2.01 0.039 0.9556 0.9346 31.56 42.12
CAD
2.25 0.14 0.9119 0.7920 19.74 30.89
CAG
1.43 0.025 0.7045 0.4308 34.97 43.84
Table 20: The parameters of intraparticle diffusion model for adsorption of Fe3+ (pH=2.6) onto CA and CACs.
Intrapartical diffusion modle
Ki1
(mg/g m0.5)
Ki2
(mg/g m0.5)
Ri12
Ri22
Intercept i1
Intercept i2
CA
2.75 0.097 0.9892 0.108 33.98 46.16
CAB
3.79 0.11 0.7902 0.1968 27.28 43.7
CAMK
1.63 0.012 0.9593 0.9589 36.36 44.85
CAHA(I)
0.3 0.0017 0.6592 0.3204 48.14 49.65
CAHA(II)
0.17 0.013 0.9673 0.5278 48.37 49.25
CAD
1.33 0.020 0.9937 0.9134 35.9 43.24
CAG
0.62 0.062 0.9351 0.7898 41.65 45.22
from Tables (18-20) it can be conclouded that for adsorption of Pb2+(pH=5.7), Pb2+ (pH=4) and Fe3+ (pH=2.6) onto CA and CACs, the calculated values of ki1 were higher than that of ki2. The reason could be as circumscription of the available vacant space for diffusion in them, so of pore blockage. The values of the correlation regression coefficients characterizing the applicability of the intraparticle diffusion model (Ri12, Ri22) were lower than that of R2 (pseudo second order), but commensurable with R12 ( pseudo first order). Actually, the three models stated above could describe the proposed sorption process to a definite extend, but they could not predict the high rate of adsorption during the first minutes of the process. Probably, the initial stages are controlled by external mass transfer or surface diffusion, followed by chemical reaction or a constant-rate stage, and diffusion causing gradual decrease of the process rate [192]
3.5. Adsorption isotherms:
Adsorption isotherm studies are necessary for illustrating the adsorption process at equilibrium conditions. An adsorption isotherm is characterized by certain constants which express the adsorbent affinity and can also be used for finding the adsorption capacity of the sorbent. The adsorption of iron and lead from polluted water using CA and CACs could be assumed to have abehavior fitting with the isothermal adsorption model in which the adsorbate keeps a dynamic equilibrium between the adsorotion and desorption at afixed temperature [178, 200].
Two most widely used mathematical models Langmuir and Freundlich adsorption isotherms are adopted for expressing the quantitative relationship between the extent of sorption and the residual solute concentration. Langmuir adsorption isotherm assumes monolayer coverage of adsorabate over ahomogeneous adsorbent surface and the adsorption of each molecule onto the surface has the same activation energy of adsorption.
Freundlich adsorption isotherms assumes aheterogeneous surface with anon-uniform distribution of heat of adsorption over the surface with the possibility of the multilayer adsorption [178, 179]. The maximum metal ions adsorption capacities are determined by analyzing the experimental data for heavy metal adsorption onto CA and CAHA(I) [201], as they provide the higher removal effeciency. The data of Pb2+ and Fe3+ adsorption by CA and CAHA(I) are examined in accordance with langmiur adsorption isotherm models whose linearized equation was:
Ce/qe =Ce/qm + 1/ KL qm
where qm (mg /g) and KL (L /mg) are constants in Langmuir’s equation which are referred to the maximum adsorption capacity corresponding to complete monolayer coverage and the Langmuir model constant that is indirectly related to the energy of adsorption. Also qe and Ce parameters represent the equilibrium adsorption capacity and the equilibrium concentration heavy metal ion that is remaining in solution, respectively. qe is calculated as follows:
qe =((Co-Ce)V)/m
where Co is the initial metal ion concentration (mg/L), Ct is the equilibrium concentration of adsorbate (mg/L) (mg/L), V is the initial solution volume (L) and m is the adsorbate dose (g). A plot of Ce/qe vs. Ce (Fig. 58-61) gives a linear trace with a slope of 1/qm and intercept of 1/ KL qm. A further analysis of the Langmuir equation can be made on the basis of a dimensionless equilibrium parameter, RL, also known as the separation factor,
given by
RL=1\(1+ KLCe)
where Ce is equilibrium liquid phase concentration of the solute at which adsorption is carried out. The value of RL lies between 0 and 1 for favorable adsorption, while RL > 1 represents unfavorable adsorption, and RL = 1 represents linear adsorption, while the adsorption process is irreversible if RL = 0 [179, 202].
Also the obtained data are examined in accordance with the Freundlich equation which is expressed as
qe =log KFCe1/n
where KF is the measure of adsorption capacity and n is the adsorption intensity linear form of Freundlich
log qe =log KF + (1/n) log Ce
where qe is the amount adsorbed (mg/g), Ce is the equilibrium concentration of adsorbate (mg/L), and KF and n are the Freundlich constants related to the adsorption capacity and adsorption intensity, respectively. A plot of log qe vs. log Ce (Fig. 58-61) gives a linear trace with a slope of 1/n and intercept of log KF. The 1/n value in the range of 0 and 1 is a predicting of adsorption intensity of metal ion onto the adsorbent and the type of isotherm to be irreversible (1/n=0), favourable (0<1/n<1) and unfavourable (1/n >1) [197], the 1/n value determine the surface heterogeneity, becoming more heterogeneous as its value gets closer to zero. In addition, the value of n varies with the heterogeneity of the adsorbent, if n < 10 and n > 1 indicating the adsorption process is favorable [193, 203].
The isotherm parameters and correlation coefficients calculated for the adsorption of Pb2+ and Fe3+ using CA and CAHA(I) are listed in Tables (21,22).
Table 21: Isotherm parameters and correlation coefficients calculated for the adsorption of Pb2+ using CA and CAHA(I).
Adsorbent
Langmuir Isotherm
Freundlich Isotherm
qmax(mg/g) KL R2 KF 1/n N R2
CA
51.78
1.264
0.9353
18.62
0.377
2.64
0.6028
CAHA(I)
52.99
0.702
0.8652
16.89
0.3450
2.89
0.9555
Table 22: Isotherm parameters and correlation coefficients calculated for the adsorption of Fe3+ using CA and CAHA(I).
Adsorbent
Langmuir Isotherm
Freundlich Isotherm
qmax(mg/g) KL R2 KF 1/n N R2
CA 66.53 0.0816 0.9361 15.84 0.2602 3.84 0.9656
CAHA(I) 113.63 0.1560 0.7675 24.35 0.3334 2.99 0.9178
from Table 21, it can be concluded that for adsorption of lead on CA, the Langmuir isotherm (R2 > 0.93) fitted the experimental results better than those of the Freundlich isotherm (R2 > 0.60) as reflected with the correlation coefficient, indicating the homogenous feature presented on the CA surface and demonstrates the formation of monolayer coverage of the lead ions on the CA surface, the adsorption is localized, all active sites of surface have similar energies and no interaction between adsorped molecules. Moreover, the value of RL was 0.0003. This also suggests an irreversible adsorption between CA and Pb2+ ions [202].
On the other hand, for CAHA, it can be stated that the Freundlich isotherm (R2 > 0.95) fitted the experimental results comparable to the Langmuir isotherm (R2 > 0.86), indicating that the adsorbed amount increased with initial concentration. The slope 1/n provides information about surface heterogeneity and surface affinity for the solute. As a higher value of 1/n (0.34) is obtained, it corresponds to the greater heterogeneity of the adsorbent surface. Furthermore, the value of n > 1 obtained from the Freundlich isotherm indicating (2.8), that this process is also favorable [203] and heterogeneous sorption. The maximum adsorption capacities of the Pb2+ ions are found to be 51.78 and 52.99 mg/g for CA and CAHA(I), respectively.
from Table 22, it can be stated that for adsorption of iron on CA and CAHA(I), the Freundlich isotherm (R2 > 0.96), (R2 > 0.91) fitted the experimental results comparable to the Langmuir isotherm (R2 > 0.93), (R2 > 0.76) for CA and CAHA(I) respectively. The slope 1/n provides information about surface heterogeneity and surface affinity for the solute. As a higher value of 1/n is obtained, it corresponds to the greater heterogeneity of the adsorbent surface. Furthermore, the value of n > 1 obtained from the Freundlich isotherm (3.8, 2.9), indicating that this process is also favorable and heterogeneous sorption.
Firure 58: Langmuir and Freundlish isotherms for adsorption of Pb2+ on CA Powder.
Firure 59: Langmuir and Freundlish isotherms for adsorption of Pb2+ on CAHA(I) Powder.
Firure 60: Langmuir and Freundlish isotherms for adsorption of Fe3+ on CA Powder.
.
Firure 61: Langmuir and Freundlish isotherms for adsorption of Fe3+ on CAHA(I) Powder
SUMMARY AND CONCLUSION
The present thesis comprises of there chapters:
Chapter (1) includes introduction and literature review that focused on the heavy metals found in industrial effluents as hazardous pollutants that may affect hardly the environment.
This chapter also includes various technologies which have been used to remove metal ions. Especial attention is given to two types of heavy metals ions (Pb2+ and Fe3+), their abundance, use in several industries and harmful effect on the environment especially to aquatic systems.
The literature review also focused on calcium alginate and its biomedical applications, also focused on bentonite, metakaolin, diatomeous silica, gelatin and hydroxy apatite (produced from biowastes) which form composites with calcium alginate and their role in heavy metals removal.
Chapter (2) includes experimental and methods which focused on types of chemical used, methods of preparation of calcium alginate, hydroxy apatite and calcium alginate composites and characterization tecniques such as X-Ray Diffraction (XRD), Fourier Transformer Infrared (FTIR), Scanning Electron Microscope(SEM) and Energy Dispersive Analysis X-ray (EDX). Preparation of heavy metal solutions with the assessment of their concentrations before and after adsorption process using Perkin Elmer 2380 Atomic Absorption Spectro Photometer. The six synthesized composite beside calcium alginate used in the present study are listed in the following table:
Composite Compound Source of raw material Abbrev.
1 Calcium alginate Oxford Lab. Reagent CA
2 Calcium alginate-Bentonite Bentonite (Abu Zaabal Fertilizer & Chemicals Co. CAB
3 Calcium alginate-Metakaolin Metakaolin (Sinai Peninsula) CAMK
4 Calcium alginate-HAP(1) HAP (egg shell calcined at 900oC) CAHA(I)
5 Calcium alginate-HAP(2) HAP (bovin bone calcined at 1000oC) CAHA(II)
6 Calcium alginate-Diatomeous Diatomeous (Kazakhstan) CAD
7 Calcium alginate-Gelatin Oxford Lab. Reagent CAG
Chapter (3) is concerned with the results and discussion that includes characterization of the prepared composites before and after heavy metal uptake using XRD, FTIR, SEM and EDX techniques. It also include the adsorption process of heavy metal uptake and kinetics studies of the adsorption process.
I. Fourier Transform Infrared (FTIR) analysis:
The FTIR spectra of calcium alginate showed important absorption bands regarding hydroxyl, ether and carboxylic functional groups. Stretching vibrations of O–H bonds of alginate appeared at 3444 cm-1. In the spectra of CAG, the absorption band at around 3442 cm-1 concerned with OH stretching vibration for CA slightly broadened and shifted to a lower wave number with the blending with gelatin, suggesting the formation of an intermolecular hydrogen bond.
CA raw material, reveals asimilar spectra with CAB. A strong and sharp band is detected at 1022 cm-1 which is related to Si–OH stretching vibrations. 1034 cm-1 assigned to Si-O stretching, and 875 cm-1 is due to OH bending of the Al-Al-OH group. A similar OH bending vibration is observed for Al-Mg-OH at 842 cm-1, 690 cm-1 assigned to quartz. Also, there is a shoulder peak at 520 cm-1 (Al-O-Si bending), and 464 cm-1 (Si-O-Si bending). The band of Si-O-Si bending in CAD is the strongest and sharpest compared to CAMK and CAB.
The bands corresponding to the samples of CAHA(I) and CAHA(II) which appears at 3570 and 630-633 cm-1 (due to structural OH of hydroxy apatite) confirm the hydroxyapatite. The most intensive bands in the range of 1044 –1090 cm-1 corresponded to the triply degenerated asymmetric stretching vibrations of P-O. Otherwise the peak at 962.97 cm-1 indicates the non-degenerated asymmetric mode of PO43-. The very strong and sharp bands observed at 569-572 and 602-603 cm-1 attributed to triply degenerated bending mode of the O-P-O in PO43- group. The larges parting distance of these bands revels the crystalline phase.
II. X-ray diffraction for crystal phase detection:
The X-ray diffraction pattern of CA and CACs in the range of 2θ = 5-60o diffraction degree shows two typical peaks in 2θ =16° and 22° corresponding to calcium alginate. The XRD of CAG shows typical peaks around 12° and 21°. The samples of CAB, CAMK and CAD revealed very similar XRD pattern peaks of quartz phase.
The sample CAHA(I) and CAHA(II) revealed very similar XRD pattern peaks of 100 % HAP (reference code :01-086-1194) in the rang of 2θ =31o and 32o and d (Ao) value spacing equal 2.81 and 2.78.
XRD analysis of CAHA(II) also indicated the absence of secondary phases, such as tri calcium phosphate (TCP) or calcium oxide (CaO). In the case of CAHA(I), its diffraction pattern revealed additional phase of β-tricalcium phosphate, beside the HAP as a main phase. Hydoxyapatite prepared from bovine bone contain certain amount of carbonate (CO32-) in its lattice structure. The carbonate ions affect on the the degree of crystallinite and hence increase the bioactivity of pure HAP.
III. Scanning Electron Microscope (SEM) and Energy Dispersive X-ray (EDX)
According to the results obtained by EDX analysis, carbon, oxygen and calcium are the major constituents of all the prepared composites related to calcium alginate polymer present. EDX analyses of CA and CAG revealed that the major elements present are carbon and oxygen corresponding to polymer composition. The SEM of CAG showed a smooth and homogeneous morphology, suggesting high miscibility and blend homogeneity between calcium alginate and gelatin.
In the case of CAB, CAMK and CAD, the presence of an obvious peak related to the Si compounds is evident. According to the results obtained by EDX analysis, weigh percent of elements present indicating a large portion of the composites is composed of Si compounds which is suitable for an efficient sequestering metal cations from aqueous solution.
EDX analyses of the prepared HAP revealed that inorganic phases of bovin bone and egg shells were mainly composed of calcium and phosphorus as the major constituents with some minor components such as C, O, Na, Mg and Si. The weight and atomic percentage shows that the Ca/P ratio around 1.7 and 1.8 which is below 2 and acceptable where the ideal Ca/P ratio of HA is 1.67, in HAP composites (CAHA(I) and CAHA(II)) this ratio increased than 2 due to the excess amount of Ca crosslinkage of calcim alginate presents.
IV. FTIR Analysis after metal ion uptake:
FTIR spectra of CA and CAHA(I) after targeted metal ions uptake ( Pb2+, Fe3+ ) showed that there is no significant change in peak positions after metal ion uptake and also no new absorption peaks are detected . There are alittle peak shifts which may be attributed to the corporation and substitution of metal ion in the lattice structure of CA and CAHA(I).
V. X-ray diffraction for crystal phase detection after metal ion uptake:
XRD patterns of CA and CAHA(I) after Pb2+ and Fe3+ metal ions removal did not revealed any new phases. supported the proposal that Pb2+ and Fe3+ ions uptake was not dependent on dissolution/precipitation mechanisms. Pb2+ and Fe3+ ions removal may be occurs by adsorption mechanisms like surface complexation or ionic exchange. The XRD patterns showed some changes in their relative intensities and crystal sizes. Also, ther are some little shifts in d- spacing values, this may be due to the ion exchange between Ca2+ and metal cations of Pb2+ and Fe3+ in lattice structure of CA and CAHA(I) Powder.
VI. Scanning Electron Microscope (SEM) and Energy Dispersive X-ray (EDX) after metal ion uptake:
SEM of CA and CAHA(I) after metal ion uptake revealed some changes in morphology and microstructure. Morever EDX results indicated the presence of Pb2+ and Fe3+ ions with CA and CAHA(I). the peak of Fe3+ is more intense than that of Pb2+ and this confirms, the higher value of Fe3+ ions uptake by CA and CAHA(I) comparing with Pb2+ ions.
The weight and atomic percentage of Ca2+ ions in CA and CAHA(I) after metal ion uptake was less than its value before metal ion removal. The decrease in calcium percentage may be attributed to ion exchange between targeted metal ions and; (1)calcium ions of calcium alginate in CA and CAHA(I), (2) or Ca ions of HA present in CAHA(I). This ion exchange mechanism between Pb2+ ions (as example) and Ca2+ ions of HA produced anew phase of hydroxypyromorphite. However , this phase is not detected in the present study this may be attributed to under limit of XRD.
Also the uptake process may be occurs during chelation bonding of targeted metal ions with two carboxylic groups of alginate and one or two OH sites of the alginate. In this case metal ions may forms complexes with two adjacent alginate rings. Here,‘‘adjacent’’ means either two neighbor alginate rings of a single polymeric chain (intramolecular chelation) or two rings from two parallel chains (intermolecular chelation).
VII. Metal ion uptake by adsorption process.
1- Effect of contact time on the adsorption process.
The effect of contact time on the adsorption capacity of CA and CACs for Pb2+ ( natural pH of 5.7), Pb2+ (pH=4) and Fe3+ ( natural pH of 2.6) indicated that, the equilibrium points of adsorption are attained within the first 10 – 60 min. of contact time with different removal efficiency and slightly similar exponential phase as listed in table below. the adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+(pH=2.6)onto CA and CACs powder was rapid in the initial stages and was almost same at high contact time. This behavior could be attributed to the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption process. Further increase in contact time led to no significant adsorption of metal ions by the adsorbents probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. As it can be seen in the following table, CAG is the higher efficiency for lead (Pb2+) removal then CAHA(I) and CA. On the other hand, CAHA(I) is the higher efficiency for iron (Fe3+) removal then CAHA(II) and CA.
Equilibrium point time (min.) Removal effeciency(%) Exponential phase time (min.)
Pb2+
(pH=5.7)
Pb2+
(pH=4)
Fe3+
(pH=2.6)
Pb2+
(pH=5.7)
Pb2+
(pH =4)
Fe3+
(pH=2.6)
Pb2+
(pH=5.7)
Pb2+
(pH=4)
Fe3+
(pH=2.6)
CA
60
30
30
82.3
86.82
94.15
5-60
5-30
5-30
CAB
60
30
30
66.4
68.2
89.42
5-60
5-30
5-30
CAMK
30
30
30
69.26
77.7
89.87
5-30
5-30
5-30
CAHA(I)
60
30
30
80.78
86.62
99.33
5-60
5-30
5-30
CAHA(II)
30
30
10
74.08
84.67
98.65
5-3
5-30
5-10
CAD
30
30
30
68.96
63.23
86.72
5-30
5-30
5-30
CAG
30
30
30
83.43
87.93
91.11
5-30
5-30
5-30
2- Effect of pH
The effect of pH on the adsorption is performed only for Pb2+ because of Fe3+ solution was stable only at pH lower than 3. The study is achieved with two pHs values (4 and 5.7) the original solution of lead is at pH=5.7 and pH=4 at 5-120 minutes contact time, 20 mg dosage of the different composites and 100 ppm of metal solution at 25±1°C. The adsorption efficiency of Pb2+ at pH=4 is higher than that of pH 5.7 for all the composites.
3- Effect of adsorbent dosage on metal ion adsorption.
The experimental results of the adsorption of Pb2+ on CA ( as astandard model ) as afunction of adsorbent dosage 10, 15 and 20 mg/10 mL, initial Pb2+ concentration of 100 mgL-1, natural pH of 5.7, temperature 25oC at the optimal contact time (30 min) and interval contact time ( 5-30 min) showed that, the Pb2+ adsorption percent rapidly increased with the increase in the adsorbent dosage . this can be attributed to higher adsorbent dosage due to the increased surface area providing more adsorption sites available which gave rise to higher removal of lead.
VIII. kinetics studies of the adsorption process.
The kinetic study is useful to predict the adsorption rate which is very important in modeling and designing of the adsorption process. The kinetic of adsorption are evaluated at an initial concentration of 100 mg/L for Pb2+(pH=5.7),Pb2+( pH 4)and Fe3+(pH=2.6), adsorbent dosage of 0.02 g/10 mL and temperature of 25oC. By appling the pseudo-first rate equation of lagergren, it is clear that the regression coefficient does not close to unity. Also, the values of qe obtained from pseudo-first order equation for all the adsorbent are different and not matched notably with the experimental qe value. from the linear plots of pseudo-second rate equation of lagergren, the qe,experimental and the qe,calculated values are very close to each other, and also, the calculated coefficients of determination, R2, are close to unity
from all the obtained results, it is obvious that the regression coefficient (R2) from pseudo-second order rate equation for all the adsorbents was higher than that of the pseudo-first order model. On the basis of the regression coefficient and calculated values of adsorption capacity, the adsorption process was found to obey and exhibited best fit to the pseudo-second-order kinetic model which is mean that the rate-limiting step might be chemical adsorption or chemisorption involving valency forces through exchange of electrons between the sorbate and the sorbent, also only one ion of the metal is sorbed onto two sorption sites on the sorbent surface.
IX. Prediction of adsorption rate-limiting step
There are essentially three consecutive mass transport steps associated with the adsorption of solute from the solution by an ads0rbent. These are (1) film diffusi0n, (2) intraparticle or p0re diffusion, and (3) sorption into interior sites. The third step is very rapid and hence, film and pore transports are the major steps controlling the rate of adsorption. The most commonly used technique for identifying the mechanism involved in the adsorption process is by fitting an intraparticle diffusion plot proposed by Weber and Morris.
The results stated that the sorption process proceeds by surface sorption and intraparticle diffusion. The initial rapid uptake can be attributed to the boundary layer effects (film diffusion). After the external surface loading was completed, the intraparticle diffusion or pore diffusion takes place. However, the plot indicated that the intraparticle diffusion was not the rate-controlling step because it did not pass through the origin.
X. Adsorption isotherms:
Adsorption isotherm studies are necessary for illustrating the adsorption process at equilibrium conditions. Two most widely used mathematical models Langmuir and Freundlich adsorption. Langmuir adsorption isotherm assumes monolayer coverage of adsorabate over ahomogeneous adsorbent surface and the adsorption of each molecule onto the surface has the same activation energy of adsorption. Freundlich adsorption isotherms assumes aheterogeneous surface with anon-uniform distribution of heat of adsorption over the surface with the possibility of the multilayer adsorption
The results of single metal ion adsorption of Pb2+ onto CA at 25 oC can be represented well by langmiur than Freundlish model with good correlation coefficient (R2). This means that the adsorbates containing Pb2+ was adsorbed in such amanner that only one atomic layer of adsorbate can be adsorbed and distributed uniformly on the surface of the adsorbents (CA) and the adsorption of each molecule onto the surface has the same activation energy of adsorption. the value of RL was 0.0003. This also suggests an irreversible adsorption between CA and Pb2+ ions.
For iron and lead adsorbed onto CAHA and iron adsorbed into CA, it can be stated that the Freundlich isotherm well fitted the experimental results comparable to the Langmuir isotherm indicating that the adsorbed amount increased with initial concentration. The slope 1/n provides information about surface heterogeneity and surface affinity for the solute. As a higher value of 1/n is obtained, it corresponds to the greater heterogeneity of the adsorbent surface. Furthermore, the value of 1 < 1/n > 0 and the value of n > 1 obtained from the Freundlich isotherm indicating, that this process is also favorable and heterogeneous sorption.
from all the obtained results and analysis we can stated that, the uptake of Pb2+ and Fe3+ may be occurs by adsorption mechanisms like surface complexation during chelation bonding of targeted metal ions with two carboxylic groups of alginate and one or two OH sites of the alginate ring forming complexes with two adjacent alginate rings. Here,‘‘adjacent’’ means either two neighbor alginate rings of a single polymeric chain (intramolecular chelation) or two rings from two parallel chains (intermolecular chelation). or ion exchange between targeted metal ions and; (1) Calcium ions of calcium alginate in CA and CACs, (2) or Ca ions of HA present in CAHA(I) and CAHA (II).
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الملخص العربي
الرساله تنقسم الي ثلاثة فصول رئيسية :
الفصل الاول : ويتضمن المقدمة والمراجع التاريخية التي تخص العناصر الثقيلة كمُلوثات بيئية والتي تنتج بكمية كبيرة من المخلفات الصناعية والتي تؤثر بصورة خطيرة علي البيئة وصحة الانسان. وشمل هذا الفصل أيضا مختلف التقنيات المستخدمه لأزالة هذه العناصر. وتم القاء الضوء علي عنصرين من هذه العناصر وهما الرصاص والحديد من حيث مصادر تواجدها من المخلفات الصناعية وتأثيرها الضارعلي البيئة المائية. ويتضمن هذا الفصل أيضا شرح وافي لألجينات الكالسيوم وتطبيقاتها في ازالة العناصر الثقيلة, والجيلاتين وبعض الطفلات مثل البنتونيت والميتاكاولين وسيليكا الدايتموس والهيدروكسي اباتيت المحضرة من المخلفات الحيوية (مثل قشر البيض وعظام البقر) والتي تكون مركبات مع الجينات الكالسيوم للتخلص من العناصر الثقيلة.
الفصل الثاني : ويشمل الكيماويات المستخدمة , وطرق تحضير ألجينات الكالسيوم والهيدروكسي اباتيت ومركبات ألجينات الكالسيوم المختلفة وطرق التعرف عليها باستخدام قياسات حيود الاشعة السينية , الأشعة تحت الحمراء, الميكرسكوب الألكتروني الماسح للضوء وطاقة الاشعة السينية المشتتة. وكذلك طرق تحضير محاليل ايونات العناصر الثقيلة وتقديرتركيزها قبل وبعد عملية الإمتزاز باستخدام جهاز الامتصاص الذري. ويتضمن الجدول التالي مصادر المخلفات الحيوية والمواد الخام وألجينات الكالسيوم ومركباته السته المحضرة :
Composite Compound Source of raw material Abbrev.
1 Calcium alginate Oxford Lab. Reagent CA
2 Calcium alginate-Bentonite Bentonite (Abu Zaabal Fertilizer & Chemicals Co. CAB
3 Calcium alginate-Metakaolin Metakaolin (kaolin (Sinai Peninsula) calcined at 800oC CAMK
4 Calcium alginate-HAP(1) HAP (egg shell calcined at 900oC) CAHA(I)
5 Calcium alginate-HAP(2) HAP (bovin bone calcined at 1000oC) CAHA(II)
6 Calcium alginate-Diatomeous Diatomeous (Kazakhstan) CAD
7 Calcium alginate-Gelatin Oxford Lab. Reagent CAG
الفصل الثالث : ويتناول النتائج التي تم الحصول عليها ومناقشتها من حيث توصيفها قبل وبعد عملية الامتزاز باستخدام اجهزة SEM وEDX مع XRD و FTIR ويشمل ايضا دراسة عملية الامتزاز و حركية عملية الامتزاز.
التعرف علي المجموعات الوظيفية من قياسات الاشعة تحت الحمراء(i)
اظهرت تحاليل عينة الجينات الكالسيوم تواجد حزم امتصاص هامة تعود الي المجموعات الوظيفيه الخاصة ب الهيدروكسيل والكربوكسيليك والايثر. حيث تظهر مجموعة الهيدروكسيل الخاصه ألجينات الكالسيوم CAعند طول موجي حوالي 3444 cm-1 , ويظهر طيف مركب الجينات الكالسيوم مع الجيلاتين CAG أن ذروة الامتصاص عند حواليcm-1 3442 والخاصة باهتزاز مجموعة الهيدروكسيل OH لـ ألجينات الكالسيوم قد اتسعت قليلاً وتحركت إلى طول موجي أقل بالمزج مع الجيلاتين ، مما يشير إلى تكوين رابطة هيدروجينية بين جزيئية.
يوجد تشابه كبير بين طيف ألجينات الكالسيوم وطيف ألجينات الكالسيوم مع البنتونيت CAB والميتاكاولين CAMK وسيليكا الدايتاموس CAD , ففي CAB تظهر ذروة امتصاص حادة وقوية عندcm-1 1022 والخاصة باهتزاز مجموعة Si-OHوعند cm-1 1034 لمجموعة Si-O وعند 875 cm-1 بسبب وجود OH bending لمجموعةAl-Al-OH كذلك تهتز مجموعة مشابهه منOH bending , خاصة ب Al-Mg-OH تظهر عند842 cm-1 و cm-1 690 تعود الي تواجد الكوارتز. ايضا يوجد حزمة امتصاص شولدر عند 520 cm-1 تعود الي (Al-O-Si bending ) وعند 464 cm-1 خاصة ب (Si-O-Si bending). تلاحظ ايضا ان حزمة الامتصاص الخاصة ب Si-O-Si bending في عينة مركب ألجينات الكالسيوم مع الداياتاموس تكون اقوي وأحد مقارنة بعينات CAB و CAMK.
بالنسبة لعينات ألجينات الكالسيوم مع الهيدروكسي أباتيت المحضرة من قشور البيض وعظام البقر CAHA(I) و CAHA(II) فتظهر حزم امتصاص عند 3570 cm-1 و 630-633 cm-1 بسبب اهتزازمجموعة OH للهيدروكسي أباتيت. كما أظهرت النتائج ايضا أن ذروة الامتصاص الاكثر شده في المدي من 1044 cm-1 الي 1090 cm-1 ترجع الي اهتزازات الرابطة P-O في مجموعة الفوسفات وتكون تللك الاهتزازات متماثلة , اما عند 962.97 cm-1 فتكون تلك الاهتزازات لمجموعة الفوسفات غير متماثلة. حزم الامتصاص القوية جدا والحادة والتي تظهر عند 569-572 cm-1 و عند 602-603 cm-1 فتعزي الي اهتزازات مجموعات O-P-O في مجموعة الفوسفات PO43-.
(ii) قياسات حيود الاشعه السينية :
بينت التحاليل باستخدام X-ray diffraction لعينات ألجينات الكالسيوم ومركباته في في المدي 5-60o=θ2 تواجد عدد ذروتين خاصة بالجينات الكالسيوم في نطاق =16 o , 22 oθ2 واظهرت النتائج ايضا تواجد الجيلاتين في عينة CAG في نطاق =12 o , 21 oθ2 وتظهر نتائج تحاليل عينات CAB, CAD, CAMK تشابه كبيربينهم و تكون الطور الخاص بالكوارتز.
كما تظهر النتائج تشابه كبير لعينات الجينات الكالسيوم مع الهيدروكسي اباتيت المحضر من قشور البيض CAHA(I) و CAHA(II) ذات الهيدروكسي اباتيت المحضر من عظام الابقار وتكون طور الهيدروكسي اباتيت طبقا للكود المرجعي 01-086-1194 في نطاق =31 o , 32 oθ2 عند قيم مسافات تباعد (d-spacing) مساوية 2.81 و 2.78انجستروم , ولا تحتوى عينات CAHA(II) على أي أطوار اخرى مثل أكسيد الكالسيوم CaO أوفوسفات الكالسيوم في حين ان عينات CAHA(I) تظهر تواجد طور فوسفات الكالسيوم بجانب طور الهيدروكسي اباتيت كطور رئيسي وتظهر نتائج تحاليل عينة الجينات الكالسيوم مع الهيدروكسي اباتيت المحضرة من عظام الابقار CAHA(II) انها تحتوي علي كمية من الكربونات في الشبكة البلورية. وتوفر ايونات الكربونات زيادة في النشاطية الحيوية للهيدروكسي أباتيت كما تؤثر في درجة التبلور للمركب.
(iii) التعرف على الشكل المورفولوجى والتحليل النوعي والكمى للعينات المحضرة بأستخدام قياسات الميكروسكوب الإلكترونى الماسح للضوء وطاقة الأشعة السينية المشتتة.
أظهرت صور الميكروسكوب الإلكتروني لجميع العينات تكتل بلوري للجزيئات مع شكل غير منتظم
نسبيا ، مع أحجام مختلفة من البللورات وتوضح الصور الشكل المورفولوجي المتجانس والناعم لعينة الجينات الكالسيوم مع الجيلاتين مشيرة الي الخلط المتجانس بينهما. وأثبتت خرائط EDX لجميع العينات تواجد اشعاع Kα لعناصر الكربون والاكسجين والكالسيوم والمفترض تواجدها في بوليمرات ألجينات الكالسيوم وتظهر النتائج ايضا في العينات المحتوية علي الهيدروكسي أباتيت ان الأطوار غير العضوية في قشور البيض وعظام الابقار تتكون اساسا من عناصر الكالسيوم والفوسفور بالاضافة لكميات قليلة من عناصر الكربون والاكسجين والصوديوم والماغنسيوم بالاضافة الي السيليكون. مع الحصول على نسبة مولارية لعنصر الكالسيوم مع الفوسفور (P/Ca) في الهيدروكسي أباتيت المحضرة معمليا مساوية 1.67 ولكن تزيد هذه النسبة بسبب تواجد كميات زائدة من كالسيوم الروابط البينية في الجينات الكالسيوم.
بالنسبة لعينات CAB و CAMK و CAD فيوجد اشعاع لعنصر السيليكون مُثبتاً مساهمة مركبات السيليكون في هذه المخاليط بنسب كبيرة والتي تساعد في التخلص من كاتيونات العناصر الثقيلة بدرجة كبيرة.
التعرف علي المجموعات الوظيفية من قياسات الاشعة تحت الحمراء بعد عملية الامتصاص (iv)
اشارت نتائج FTIR انه لايوجد تغير كبير في مواضع حزم الامتصاص بعد امتصاص ايونات المعادن وكذلك لم يتم الكشف عن اي ذٌرو امتصاص جديدة. ومع ذلك , هناك قليل جدا من الإزاحة لذروات الامتصاص , ويمكن ان يٌعزي ذلك الي اشتراك واحلال المعادن في الشبكة البلورية لألجينات الكالسيوم ومخاليطة المختلفة.
(v) قياسات حيود الاشعه السينية لعيينات للمركبات المحضرة بعد عملية الامتصاص
أجريت تحاليل الاشعة الشعة السينية XRD لعينات CA وCAHA(I) باعتبارهما الاعلي في عملية الامتصاص. تظهر النتائج انه لايوجد اطوار جديدة بعد امتصاص ايوني الرصاص الثنائي والحديد الثلاثي مما يدعم افتراض عدم حدوث الامتزاز نتيجة ميكانيكية التفكك والترسيب وانه ربما تحدث نتيجة عملية الامتزاز مثل التبادل الايوني وتكوين المتراكبات , وتظهر النتائج ايضاً حدوث تغيرات طفيفة في الشدة النسبية والحجوم البلورية وكذلك قيم التباعد d- spacing ويمكن ان يعزي ذلك الي التبادل الايوني بين ايونات المعادن و (1) ايونات الكالسيوم الموجودة في ألجينات الكالسيوم في CA و CAHA(I) كما في التفاعل التالي :
Ca(ALG)2 + Pb2+ Pb(ALG)2 + Ca2+
أو (2) ايونات الكالسيوم الخاصة بالهيدروكسي اباتيت في عينة ال CAHA(I) مكونا طور جديد من الهيدروكسي بيرومورفيت والتي لم تظهر في تحاليل XRD ربما لانها تكون بنسبة ضئيلة جدا تحت المدي الحثي XRD.
Ca10 (PO4)6 (OH)2 + x Pb2+ x Ca2+ + Ca10-xPbx (PO4)6(OH)2
(vi) التعرف على الشكل المورفولوجى والتحليل النوعي والكمى للعينات المحضرة باستخدام قياسات الميكروسكوب الإلكترونى الماسح للضوء وطاقة الأشعة السينية المشتتة بعد امتصاص ايونات الفلز.
كشفت النتائج عن بعض التغيرات في الشكل والبنية المجهرية ل CA و CAHA(I) عند التفاعل مع ايونات معادن الحديد والرصاص علاوة علي ذلك اشارت نتائج EDX الي وجود ايونات Pb2+ وFe3+ , كما تظهر النتائج أن ايونات الحديد الثلاثي هي اكثر كثافة من ايونات الرصاص الثنائي وهذا يوكد علي ان ايونات Fe3+ اكثر ازالة بواسطة CA و CAHA(I) مقارنة بايونات Pb2+. وتظهر النتائج ايضا انخفاض في نسبة ايونات الكالسيوم بعد عملية الامتصاص عنها قبل عملية الامتصاص, وهذا يمكن إيعازه الي احتمالية حدوث عملية الامتصاص نتيجة التبادل الايوني والذي يتوافق مع نتائج قياسات FTIR و تحاليل XRD. أيضا ربما تحدث عملية الامتصاص نتيجة ارتباط مخلبي لأيونات الفلزات مع مجموعتين كربوكسيليتين من الألجينات ومجموعة أواثنين من مجموعات الهيدروكسيل للألجينات , وفي هذه الحالة ربما يُكًون أيون الفلز متراكبات مع حلقتين متجاورتين لسلسلة بوليمرية واحدة من الألجينات او حلقتين لسلسلتين متوازيتين.
(vii) دراسة معدل امتصاص ايونات الفلزات تحت تاثير زمن التلامس
تمت دراسة تأثير الزمن علي قدرة ألجينات الكالسيوم ومخاليطة المختلفة في التخلص من ايونات الرصاص عند إس هيدروجيني للمحلول المحضر وهو pH=5.7 و عند 4pH = وايونات الحديد عند الاس الهيدروجيني للمحلول المحضر وهو 2.6 , وتشير الدراسة الي انه تم تحقيق نقاط الاتزان لجميع المخاليط في خلال من 10 الي 60 دقيقة من بداية الامتزاز ومراحل امتزاز (5-30) دقيقة , بنسب ازالة تتراوح ما بين66% إلي 83% للرصاص عند pH=5.7 و من 68% إلي 87% عند pH=4 وتتراوح مابين86% إلي 99% للحديد عند pH=2.6 .
وتكون عملية الامتزاز سريعة في المراحل الاولي ومتساوية تقريبا عند ازمنة التلامس العالية ويٌعزي هذا السلوك الي توافر عدد كبير من المواقع النشطة خلال المراحل الاولي من عملية الامتزاز وبالزيادة الكبيرة في زمن التلامس لم يُلاحظ حدوث أي عملية امتزاز وذلك بسبب الانخفاض في معدل الانتشار حيت ان جميع المواقع قد تم تغطيتها بايونات الفلز.
(viii) دراسة معدل امتصاص ايونات الفلزات تحت تاثير درجة الحموضة
تم دراسة تغيير درجة الإس الهيدروجيني علي عملية الامتزاز بالنسبة لايونات الرصاص فقط حيت ان ايونات الحديد Fe3+ تكون ثابته فقط عند اس هيدروجيني اقل من 3. تمت الدراسة عند قيمتين للإس الهيدروجيني وهما 5.7 (pH للمحلول المحضر) وعند pH=4, تركيز 100 ملجرام/لتر من ايونات الرصاص وكتلة من المادة المازّة ( (adsorbent 20 ملجرام/ 10 ملليتر من محلول ايون الفلز عند درجة حرارة 25 درجة مئوية. اظهرت الدراسة ان كفاءة عملية الإمتزاز عند 4 = pH تكون أعلي منها عند 5.7 = pH لجميع المخاليط.
(ix) دراسة معدل امتصاص ايونات الفلزات تحت تاثير كتلة الممتز
تم دراسة تاثير كتلة المادة المازّة بالنسبة لمعدل الامتصاص وذلك بالنسبة لإمتزاز الرصاص علي مركب ألجينات الكالسيوم (كنموذج قياسي لبقية المخاليط). تمت الدراسة باستخدام كتل مختلفة 10 و 15 و 20 ملجرام من المادة المازّة و10مليلترمن محلول الرصاص بتركيز 100 ملجرام/لتر و5.7 = pH ودرجة حرارة 25 درجة مئوية وازمنة تلامس من 5 الي 30 دقيقة. وتظهر النتائج ان معدل الامتصاص يزداد بزيادة كتلة المادة المازّة وهذ يٌعزي الي زيادة مساحة السطح ومن ثم زيادة المواقع النشطة المتاحة لعملية الامتزاز وزيادة كفاءة الازالة.
(x) دراسة كيناتيكية لعملية الإمتزاز
تمت دراسة حركية عملية الأمتزاز لمعدل امتصاص أيونات الفلزات علي سطح ألجينات الكالسيوم ومركباته المختلفة وتمت الدراسة عند ظروف تجريبية ( 25 درجة مئوية , كتلةمن المادة المازّة تساوي 20 مليجرام لكل 10 مليلتر من محاليل ايونات Pb2+ وFe3+ بتركيز ابتدائي 100 ملجرام/لتر) , وتم اختبار نموذجين حركيتين شائعتين تحت الظروف التجريبية وهما: معادلة الرتبة الأولى الكاذبة لـ Lagergren ومعائلة الرتبة الثانية الكاذبة لتحليل معدل امتزاز أبونات الفلزات علىCA و CACs.
1- نموذج الرتبة الأولى الكاذب
بتطبيق معادلة لاجرجرين لتفاعل الرتبة الاولى الزائفة التالية :
ln (qe –qt ) = ln qe – k1 t
حيث ان K1هو ثابت معدل لاجارجرين للإمتزاز (دقيقة-1) وqe وqt هى كميات العناصر الممتزة (ملجرام/جرام) عند الإتزان وعند الزمنt . ومن العلاقة البيانية بين Log(qe – qt) والزمن t, اتضح أن قيم معامل الارتباط R2تكون بعيدة وغير مقتربه من الوحدة كما ان قيم qe الناتجة من معادلة الرتبة الاولي الكاذبة تختلف تماما عن qe التجريبية , وهذا يشير الي أن معادلة Lagergren من الدرجة الأولى غير مناسبة لوصف امتزاز أيونات الفلزات المستهدفة بواسطة CA و CACs المستخدمة.
2- نموذج الرتبة الثانية الكاذب
بتطبيق معادلة تفاعل الرتبة الثانية الكاذبة التالية
t/qt = 1/k2qe2 + t/qe
على البيانات المعملية لإمتزاز ايونات الرصاص والحديد بواسطة ألجينات الكالسيوم ومركباته , حيث أن 2k هو ثابت معدل تفاعل الرتبة الثانية الزائفة (جم / ملجم. دقيقة) وqe وqt الكمية الممتزة في وحدة الكتلة عند الإتزان وعند الزمن t ، ومن العلاقة البيائية بين t/qt و الزمن t يتضح ان qe الناتجة من العلاقة الخطية لمعادلة الرتبة الثانية الكاذبة متوافقة تماما مع qe التجريبية , ووجد ايضا أن قيم معامل الارتباط تقترب جدا من الوحدة , وهذا يشير الي أن معادلة Lagergren من الدرجة الثانية مناسبة لوصف امتزاز أيونات الفلزات المستهدفة بواسطة CA و CACs المستخدمة
3 - تحديد الخطوة المتحكمة في التفاعل
من المعروف انه يوجد ثلاثة خطوات في أي عملية إمتزاز من محلول لمادة ممتزة وهي كالتالي (1) انتشار الطبقات الحدودية او الانتشار الفيلمي (2) انتشار بين الجزيئات او ثقبي (3) امتزاز علي المواقع او الاماكن الداخلية وهذه الخطوة سريعة جدا وربما لاتلاحظ , لذا فالخطوتين الاوليتين هما الخطوتين المتحكمتين في معدل الامتزاز. من اكثر الطرق المستخدمة في معرفة ميكانيكية عملية الامتزاز هي باستخدام الرسم البياني للانتشار بين الجزيئات للعالِمين ويبر وموريس بالعلاقة التالية:
qt = Kid t0.5 + C
حيث ان C ثابت و Kid هي ثابت الانتشار بين الجزيئات (ملجرام/جرام.دقييقة0.5 ( وqt هى كمية العناصر الممتزة (ملجرام/جرام) عند الإتزان. من العلاقة البيانية بين qt و t0.5 يتضح ان خطوة الانتشار بين الجزيئات ليست الخطوة المتحكمة في عملية الامتزاز لان العلاقة الخطية لا تمر بنقطة الاصل , كما توضح النتائج ان الامتصاص السريع للايونات في المرحلة الاولي يكون نتيجة الانتشار الفيلمي او تاثير الطبقات الحدودية وبعد اكتمال الأسطح الخاجية تتجه الأيونات للأنتشار بين الجزيئات اوخلال الثقوب.
(xi) دراسة الامتزاز عند درجة حرارة ثابته
من المفترض إن يكون إمتزاز ايونات العناصر + Pb2وFe3+ من الماء بواسطة CA و CAHA(I)
له سلوك يتلائم مع نموذج الإمتزاز عند ثبات الحرارة حيث أن المادة الممتزة تحافظ على الإتزان الديناميكي بين الإمتزاز وعدم الإمتزاز عند درجة حرارة ثابتة ويمكن تمثيل هذا النموذج بإستخدام معادلة لانجمير أو فريندلش. معادلة لانجمير لحساب أقصى قيمة لإمتزاز العناصر والتي تمثل بخط مستقيم تعطى من العلاقة
Ce/qe =Ce/qm + 1/ KL qm
حيثCe هو تركيز الإتزان للعنصر الثقيل المتبقي في المحلول(ملجرام/ لتر) عندما تمتز منه كمية تساوي qe. و qe هى الكمية الممتزة عند الإتزان ( ملجرام/ جم) وqm هى سعة الإمتزاز القصوى التي ترجع الى حدوث تغطية كاملة للسطح من طبقة واحدة (ملجرام / جم) و kLهوثابت لانجمير الذي يتناسب عكسيا مع طاقة الإمتزاز (ملجرام) ويمكن حساب qe من المعادلة
qe =((Co-Ce)V)/m
حيثCo هو تركيز الأيون الإبتدائي (ملجرام/ لتر) و Ct التركيز النهائي للعنصر (ملجرام/ لتر بعد مرور فترة من الزمن t , و V هو حجم المحلول الابتدائي (لتر) و m هي كمية الجينات الكالسيوم او مركبه مع الهيدروكسي اباتيت المضافة , كذلك من العلاقة التالية :
RL=1\1+ KLCe
نستطيع معرفة ما اذا كانت عملية الامتزاز مفضلة او غير مفضلة او غير عكسية او خطية حيث ان RL هي معيار الاتزان بلا ابعاد فاذا كانت قيمة RL تساوي صفر فان عملية الامتزاز تكون غير عكسية , واذا كانت تساوي 1 فان الامتزاز يكون خطي , واذا كانت بين صفر و واحد تكون مفضلة واذا كانت اكبر من الواحد تكون غير مفضلة.
اما معادلة فريندلش والتي تفترض وجود سطح غير متجانس مع توزيع غير منتظم لحرارة الامتزاز علي السطح وان عملية الامتزاز تتم علي طبقات عديدة و تمثل تلك المعادلة بخط مستقيم و تعطي من العلاقة
log qe =log KF + (1/n) log Ce
حيث KF و n هما ثوابت فريندلش ومن خلال قيمة (1/n) نستطيع معرفة ما اذا كان الامتزاز غير انعكاسي اذا كانت تساوي صفر اما اذا كانت بين 0 و 1 فان عملية الامتزاز تكون مفضلة اما اذا كانت قيمة (1/n) اكبر من 1 تكون عملية الامتزاز غير مفضلة.
من تطبيق تلك المعادلات علي النتائج التي تم الحصول عليها من عملية الامتزاز يتضح ان امتزاز ايونات الرصاص علي الجينات الكالسيوم يمكن تمثيلها بنموزج لانجمير بمعامل ارتباط جيد مايعني تغطية ايونات الرصاص لطبقه واحدة من CA , كذلك قيمة RL=0 تشير الي ان عملية الامتزاز تكون غير انعكاسية.
اما في حالة إمتزاز ايونات الرصاص والحديد علي سطح CAHA(1) وكذلك امتزاز ايونات الحديد علي سطح CA فيتضح أنها تُمثل جيدا بنموذج فريندلش بمعامل ارتباط جيد وكذلك قيمة (1/n) تشير الي ان عملية الامتزاز تكون مفضلة.
من خلال كل ماسبق من نتائج التحاليل المختلفة ومن دراسة كيناتيكية عملية الامتزاز يتضح أن:
إزالة عناصر الحديد والرصاص ربما تتم عن طريق تكوين متراكبات علي سطح ألجينات الكالسيوم ومركباته المختلفة او تتم عن طريق التبادل الايوني بين ايونات تلك العناصر وايونات الكالسيوم في عينات ألجينات الكالسيوم ومركباته المختلفة او ايونات الكالسيوم المرتبطة بالهيدروكسي اباتيت في عينات ألجينات الكالسيوم التي تحتوي علي مادة الهيدروكسي اباتيت.
مركب CAG هوالأعلي كفاءة في أزالة عنصر الرصاصPb2+ يليه مركبات CA و CAHA(I) ومركب CAHA(I) هو الأعلي كفاءة في إزالة عنصر الحديد Fe3+ يليه CAHA(II) ثم CA.
}
رسالة مقدمة من الطالب
محمد أحمد عبده عبدالله الامير
بكالريوس علوم (كيمياء – 2009)
كجزء من متطلبات الحصول علي
درجة الماجستير في العلوم
(كيمياء)
(كيمياء غير عضوية وتحليلية)
إلى
قسم الكيمياء
كلية العلوم
جامعة قناة السويس
الإسماعيلية
(2019)
دراسات علي إزالة العناصر الثقيله من المياه الملوثه باستخدام بعض المواد العضوية وغير العضويه المركبه
لجنة الأشراف التوقيع
1- أستاذ دكتور / صبري عبد الحميد القرشي ..........................
أستاذ الكيمياء غير العضوية
قسم الكيمياء
كلية العلوم
جامعة قناة السويس
2- دكتور / أيمن عبد المؤمن محمد مصطفي ..........................
مدرس الكيمياء الفيزيائية
قسم الكيمياء
كلية العلوم
جامعة قناة السويس
3- دكتور/ عباس ممدوح عباس ..........................
مدرس الكيمياء غير العضوية والتحليلية
قسم الكيمياء
كلية العلوم
جامعة قناة السويس
وافق مجلس الكلية بتاريخ / /
كما وافق السيد الأستاذ الدكتور / نائب رئيس الجامعة بتاريخ / /
علي منح درجة الماجستيرفي العلوم للطالب
محمد أحمد عبده عبدالله الامير
عنوان ارسالة:
دراسات علي إزالة العناصر الثقيله من المياه الملوثه باستخدام بعض المواد العضوية وغير العضويه المركبه
لجنة الحكم والمناقشة التوقيع
1- أستاذ دكتور / صبري عبد الحميد القرشي ..........................
أستاذ الكيمياء غير العضوية
قسم الكيمياء
كلية العلوم
جامعة قناة السويس
2- أستاذ دكتور / عصام عبدالعزيز ابراهيم كيشار ...........................
أستاذ الكيمياء غير العضوية
قسم الكيمياء
بنات عين شمس
3- أستاذ دكتور مساعد / خلود محمد ابو النور ...........................
أستاذ مساعد الكيمياء التحليلية
قسم الكيمياء
كلية العلوم
جامعة قناة السويس
وكيل الكلية لشئون الدراسات العليا عميد الكلية
أ.د/ علاء الدين عبدالعزيز سلام أ.د/ محمد سعد زغلول
Studies on Heavy Metals Removal from Polluted Water Using Some Organic and Inorganic Composite Materials
A Thesis Submitted by
Mohamed Ahmed Abdo Abd Allah Elamir
B.Sc in Chemistry
In Partial Fulfillment of the Requirements For
The Degree of master of Science (M.Sc.)
in
Chemistry (Inorganic and Analytical Chemistry)
to
Chemistry Departement
Faculty of Sience
Suez Canal University
Ismailia
(2019)
1. Introduction and Literature review
Heavy metals are metallic elements which have a comparatively high density compared to water. With the assumption that overweight and toxicity are interrelated, heavy metals also include metalloids, such as arsenic, that are able to motivate toxicity at low exposure level. Nowadays, the environmental pollution by these metals causes extreme ecological and global public health worry. Also, human exposure has risen intensely owing to an exponential increase of their use in various industrial, domestic, agricultural, and technological applications. Generally heavy metals sources involve industrial, geogenic, agricultural, pharmaceutical, domestic waste, and atmospheric sources. Environmental pollution is very notable in point source areas like mining, foundries and metal extraction, and other metal-based industrial operations [1-3].
Even though heavy metals are naturally occurring elements that are present throughout the earth’s crust, most environmental pollution and human exposure result from anthropogenic activities such as mining and smelting operations, industrial production and use, and domestic and agricultural use of metals and metal-containing compounds [3-5]. Environmental contamination can also occur through corrosion of metal, atmospheric deposition, soil erosion of metal ions and leaching of heavy metals, sediment resuspension, and metal evaporation from water resources to soil and groundwater. Natural phenomena such as weathering and volcanic eruptions have also been reported to significantly contribute to heavy metal pollution [1-3]. Industrial sources of heavy metals include metal processing in refineries, coal burning in power plants, petroleum combustion process, stations of nuclear power and high-tension lines, plastics, textiles, microelectronics, wood preservation, and paper-processing plants [6, 7]. It has been stated that metals such as iron, magnesium, manganese, chromium, molybdenum, cobalt, nickel, copper, selenium and zinc are necessary nutrients which are needed for various physiological and biochemical functions [1]. Insufficient supply of these micronutrients leads to a variety of deficiency diseases or syndromes.
Also, heavy metals are considered as trace elements owing to their presence in very small concentrations (ppb range to less than 10 ppm) in various environmental regions. Their bioavailability is affected by physical factors such as temperature, phase combination, adsorption, and sequestration. It is also influenced by chemical factors which affect speciation at thermodynamic equilibrium, kinetics of complexation, lipid solubility, and octanol/water partition coefficients. Biological factors, such as characteristics of species, trophic interactions, and biochemical/physiological adaptation, also play an important role [1, 8].
The essential heavy metals perform biochemical and physiological functions in plants and animals. They are important constituents of several key enzymes and play important roles in various redox reactions [1]. Copper, for example, serves as an essential cofactor for sundry oxidative stress-related enzymes including catalase, superoxide dismutase, peroxidase, cytochrome c oxidases, dopamine β-monooxygenase, monoamine oxidase, and ferroxidases [9-11]. Therefore, it is an essential nutrient that is incorporated into a number of metalloenzymes involved in hemoglobin formation, carbohydrate metabolism, catecholamine biosynthesis, and collagen cross-linking, elastin, and hair keratin. The ability of copper to cycle between an oxidized state, Cu(II), and reduced state, Cu(I), is used by cuproenzymes involved in reduction-oxidation reactions [9-11]. However, it is this characteristic of copper that also makes it potentially toxic as a result of the transitions between the two oxidation states Cu(II) and Cu(I) can lead to superoxide and hydroxyl radicals generation [9-12].
Additionally, too much exposure to copper causing cellular damage leading to Wilson disease in humans [11, 12]. As in the case of copper, several other essential elements are required for biologic functioning; however, an excess amount of this metals causes cellular and tissue damage leading to a variety of harmful effects and human diseases. For some elements including copper and chromium, a very small range of concentrations between useful and toxic effects is present . Other metals such as cadmium, antinomy, arsenic, aluminum, barium, beryllium, bismuth, gallium, germanium, gold, indium, lead, lithium, nickel, mercury, platinum, silver, strontium, tellurium, thallium, tin , titanium, vanadium, and uranium have no definite biological functions and are regarded as nonessential metals [1].
In biological systems, heavy metals have been reported to affect cellular organelles and components like lysosome, cell membrane, endoplasmic reticulum, nuclei, mitochondrial, and some enzymes involved in detoxification, damage repair, and metabolism [13]. Metal ions have been found to interact with cell components such as nuclear proteins and DNA, causing DNA damage and conformational changes that may lead to cell-cycle modulation, carcinogenesis, or apoptosis [13, 14]. Table (1) shows the main sources, health effects and permissible limits of various toxic heavy Metals.
Table (1): Sources, health effects and permissible limits of various toxic heavy Metals according to World Health Organization(WHO) [15, 16].
Metal
Source
Potential health effect potable water limits (ppm), WHO
Copper
Zinc
Mercury
Nickel
Cadmium
Arsenic
lead
Metal finishing industry
Electroplating
Metalliferous mining
Metal finishing industry
Electroplating, Fertilizers
Metalliferous mining
Agricultural material
Manures sewage sludge
Electronics
Waste disposal
landfill leachate
Metalliferous mining
Metal finishing industry
Electrodeposition
Manures sewage sludge
Alloys and steels
Metallurgical industries
Metalliferous mining
Agricultural materials
Fertilizers, waste disposal
Landfill leachate, electronics
Electronics
Metallurgical industries
Manures sewage
Specialist alloys and steels.
waste disposal
landfill leachate
Electronics, metallurgical
Industries
Nervous system irritation followed by depression
Liver damage, Wilson disease, insomnia
Phytotoxic, depression
Anemia, lethargy
Lack of muscular coordination
Abdominal pain
Increased thirst
Poisonous
Rheumatoid arthritis
Disturbs the cholesterol
diseases of the kidneys, circulatory system and nervous system
High conc. can cause DNA damage
Eczema of hands
Human carcinogen
High phytotoxicity
Damaging fauna
Kidney damage, renal disorder
Human carcinogen
Bronchitis, emphysema
Anemia
Acute effects in children
Skin manifestations,
visceral cancers, vascular disease
Immunotoxic
Modulation of co-receptor expression
Damage the fetal brain
diseases of the kidneys
circulatory system and
nervous system
2
3
0.001
0.02
0.03
0.01
0.01
1.1. `Mechanism of heavy metals in human being.
In past decades, several studies have been carried out to investigate the mechanism of toxicity with heavy metals [15] . Toxicity and carcinogenicity of heavy metals involve many mechanistic aspects, some of which are not clearly elucidated or understood [1].
Many studies have reported that the production of reactive oxygen species (ROS) and oxidative stress play an important and a key role in the toxicity and carcinogenicity of metals for example arsenic [17-19], cadmium [20] , chromium [21, 22] , lead [23, 24], and mercury [25, 26] . For the reason of their high degree of toxicity, these previous metals are among the priority metals that are of great public health significance. They are all systemic toxicants that are known to prompt multiple organ damage, even at lower levels of exposure. As per the US Environmental Protection Agency (US EPA) and the International Agency for Research on Cancer (IARC), these elements are also categorized as either “known” or “probable” human carcinogens based on epidemiologicalds and experimental studies showing a link between exposure and cancer incidence in humans and animals.
Oxidative stress is one of the major mechanisms behind metal toxicity [27]. The formation of large amounts of reactive oxygen species, such as superoxide anion (O2.-), hydrogen peroxide (H2O2), hydroxyl radical (HO.) and singlet oxygen (1O2), has been reported to promote the induction of oxidative stress. [15, 28] In fact, various studies connect heavy metals with oxidative DNA damage since these metals may reduce the level of the main antioxidant compounds in several animal tissues by inactivating enzymes and other antioxidant molecules [29]. In humans, oxidative stress is also responsible for various diseases, including cancer, Parkinson’s disease, Alzheimer’s disease, atherosclerosis, heart failure and myocardial infarction [30, 31].
Although Heavy metal-induced toxicity and carcinogenicity involve many mechanistic aspects, some of which are not clearly elucidated or understood. Each metal is known to have unique features and physicochemical properties that confer to its specific toxicological mechanisms of action [1].
1.1.1 BIOCHEMISTRY OF TOXICITY
The heavy metals poisoning effects are because of their interference with the normal body biochemistry in the normal metabolic processes. When this metals ingested, in the acid medium of the stomach, they are converted to their stable oxidation states (Zn2+, Pb2+, Cd2+, As2+, As3+, Hg2+ and Ag+) and combine with the body’s biomolecules such as proteins and enzymes to form stable and strong chemical bonds. The equations shown in Figure (1) display their reactions during bond formation with the sulphydryl groups (-SH) of cysteine and sulphur atoms of methionine (-SCH3) [32, 33].
Figure (1): interaction between metal with proteins and enzymes.
(A) = Intramolecular bonding; (B) = Intermolecular bonding; P = Protein; E = Enzyme; M = Metal
The metal groups or the hydrogen atoms in the above case are substituted by the poisoning metal and the enzyme is thus inhibited from functioning, whereas the protein–metal compound works as a substrate and can reacts with a metabolic enzyme. In the following scheme, equation C indicates the reaction of enzymes (E) with substrates (S) in either the lock-and-key pattern or the induced-fit pattern. In both cases, a substrate fits into an enzyme in a highly specific fashion, as aresult of enzyme chirality’s, to form an enzyme–substrate complex (E-S*) as follows [33] .
(E = Enzyme; S = Substrate; P = Product; * = Activated Complex)
While at the E-S, E–S* and E-P states, an enzyme cannot accommodate any other substrate till it is freed. Occasionally, the enzymes for an entire sequence coexist together in one multi-enzyme complex consisting of three or four enzymes. The product from one enzyme reacts with a second enzyme in a chain process, with the last enzyme yielding the final product as follows:
The final product (F) goes back to react with the first enzyme thereby inhibiting further reaction since it is not the starting material for the process. Hence, the enzyme E1 becomes incapable of accommodating any other substrate until F leaves and F can only leave if the body utilizes it. If the body cannot utilize the product formed from the heavy metal – protein substrate, there will be a permanent blockage of the enzyme E1, which then cannot initiate any other bio-reaction of its function. Therefore, the metal remains embedded in the tissue, and will result in bio-dysfunctions of various gravities. Furthermore, a metal ion in the body’s metallo-enzyme can be conveniently replaced by another metal ion of similar size. Thus Cd2+ can replace Zn2+ in some dehydrogenating enzymes, leading to cadmium toxicity. In the process of inhibition, the structure of a protein molecule can be mutilated to a bio-inactive form, and in the case of an enzyme can be completely destroyed. For example, toxic As3+ occurs in herbicide, fungicides and insecticides, and can attack –SH groups in enzymes to inhibit their bioactivities as shown below in Figure (2) [32, 33].
Figure (2): interactions between arsenic and enzyme.
The most toxic forms of these metals in their ionic species are the most stable oxidation states. For example, Cd2+, Pb2+, Hg2+, Ag+ and As3+. In their most stable oxidation states, they form very stable biotoxic compounds with the body’s bio-molecules, which become difficult to be dissociated, due to their bio-stabilities, during extraction from the body by medical detoxification therapy.
1.2. Heavy metals treatment techniques:
Heavy metal uptake from inorganic effluent and industrial wastewaters can be carried out by traditional treatment processes, such as, complexation, adsorption, coagulation, ion exchange, solvent extraction, chemical precipitation, electroplating, cementation, flotation and membrane separation. All of these processes may be physical, chemical or biological as shown in Figure (3), Some of these are illustrated in Figure (4) [34] . Different methods, such as chemical precipitations, conventional adsorption [35-37] , ion exchange [38], membrane separation techniques [39] and electro-remediation techniques are used usually for industrial wastewater treatment. Precipitation is most economical and hence widely used, but many industries still use chemical procedures for treatment of effluents due to economic considerations. The efficiency of the precipitation process is extremely decrease owing to the presence of complexing agents in wastewater , and this lead to incomplete processing and production of toxic sludge. Thus several new approaches have been studied to develop low cost effective and more efficient heavy metal adsorption techniques [40].
Biosorption is considered as a user-friendly with specific affinity, low cost and simple design so it is an effective separation and purification method for heavy metals disposal from industrial wastewater and it has been widely used for this purpose [41, 42].
Figure (3): Conventional technologies for heavy metal removal.
Sorption with sorbents made of agricultural or industrial by-products are used widely for heavy metals uptake from aqueous mediums because of their abundant availability, promising physical, low cost, and, surface and chemical characteristics [43]. Those materials and methods were widely discussed meeting their advantages.
Figure (4): Some conventional methods for metal removal.
1.2.1. Physico-chemical methods
Following methods have been used by various researchers for heavy metals uptake. Physical separation process are primarily applicable to particulate forms of metals, metal-bearing particle or discrete particles . Physical separation consists of flotation, mechanical screening, gravity concentration, hydrodynamic classification, magnetic separation, attrition scrubbing, and electrostatic separation, physical separation efficiency depends on several soil characteristics such as particulate shape, particle size distribution, moisture content, humic content, clay content, density between soil matrix ,metal contaminants and heterogeneity of soil matrix, and hydrophobic properties, magnetic properties of particle surface [40, 44].
The conventional chemical processes for heavy metals disposal from waste water contain many processes such as flotation, ion exchange, adsorption, electrochemical deposition and chemical precipitation. Factors which may limit the effectiveness and applicability of the chemical process are high content of clay/silt, calcite, humic, Ca and Fe, anions, heavy metals, or high buffering capacity [45].
A. Chemical Precipitation:
Chemical precipitation is one of the most widely used for removal of heavy metal from inorganic effluent in industry because of its simple operation[46]. These conventional chemical precipitation processes yield insoluble precipitates of heavy metals as hydroxide, carbonate, phosphate and sulfide. The mechanism of this process is depend on to produce insoluble metal precipitation by reacting dissolved metals in the solution and precipitant. In the precipitation method very fine particles are generated and chemical precipitants, flocculants and coagulation processes are used to increase their particle size in order to remove them as sludge [45, 46]. As soon as the metals precipitate and form solids, they can easily be removed, and metal with low concentrations, can be discharged. Removal percentage of metal ions in the solution may be improved to optimum by changing major parameters such as pH, initial concentration, ions charge, temperature,.…etc. Hydroxide treatment is the most commonly used precipitation technique owing to its relative simplicity, low cost of precipitant (lime), and ease of automatic pH control. The solubilities of the various metal hydroxides are minimized for pH in the range (8-11).
B. Coagulation and Flocculation:
The coagulation-flocculation mechanism is based on zeta potential (ζ) measurement as the standard to define the electrostatic interaction between coagulant-flocculant agents and pollutants [47] .Coagulation process is reduced the net surface charge of the colloidal particles to stabilize by electrostatic repulsion process [40]. Flocculation process continually increases the particle size to discrete particles through additional collisions and interaction with inorganic polymers formed by the organic polymers added [48]. The minute discrete particles are flocculated into larger particles, they can be removed or separated by filtration, floatation or straining. Sludge production, transfer of toxic compounds into solid phase and application of chemicals are main disadvantages of this process.
C. Electrochemical Treatments:
Electrolysis: Electrolytic recovery is one technology used for removing metals from waste water streams. This process uses electricity to pass a current through an aqueous metal-bearing solution containing a cathode plate and an insoluble anode. Electricity can be generated by movements of electrons from one element to another. Electrochemical process to treat wastewater containing heavy metals is to precipitate the heavy metals in a weak acidic or neutralized catholyte as hydroxides. Electrochemical treatments of wastewater involve electro-deposition, electro-coagulation, electro-flotation and electro-oxidation [49].
Electro-destabilization of colloids is called coagulation and precipitation by hydroxide formation to acceptable levels. It is the most common heavy metal precipitation process forming coagulants by electrolytic oxidation and destabilizing pollutants to form folc [50]. The electro-coagulation process the coagulant is generated in situ by electrolytic oxidation of an appropriate anode material. In this process, charged ionic metal species are removed from wastewater by allowing it reacting with anion present in the effluent. This process is characterized by reduced production of sludge, ease of operation and no requirement for chemical use.
However, chemical precipitation requires a large amount of chemicals to reduce metals to permissible limit for discharge. Other drawbacks are huge sludge production, poor settling, slow metal precipitation, long-term environmental impacts of sludge disposal , and the aggregation of metal precipitates [51]. It converts the aqueous pollution problem to a solid waste disposal problem without recovering the metal.
D. Ion Exchange:
Ion exchange can attract soluble ions from the liquid phase to the solid phase. It considered is the most widely used technique in water treatment industry. As a cost-effective method, ion exchange process usually involves convenient operations and materials with low-cost, and it has been demonstrated to be very effective for elemination heavy metals from aqueous mediums, specific for treating water with low heavy metals concentration [52, 53]. In this technique cations or anions containing special ion exchanger is used to eliminate metal ions from the solution. Commonly used ion exchangers are synthetic organic ion exchange resins. It can be used only low concentrated metal solution and this method is extremely sensitive with the pH of the aqueous phase.
Ion exchange resins are water-insoluble solid substances which can absorb negatively or positively charged ions from an electrolyte solution and release other ions with the same charges into the solution in an equivalent amount. The positively charged ions in cationic resins such as sodium and hydrogen ions are replaced with positively charged ions, such as, copper, zinc and lead ions, in the solutions. In a similar way, the negative ions in the resins such as hydroxyl and chloride ions can be exchanged by the negatively charged ions such as nitrate, chromate, sulfate, cyanide, and dissolved organic carbon (DOC).
E. Membrane Filtration:
Membrane filtration has received a great attention for the remidation of inorganic effluent. It is able to remove organic compounds, suspended solid, and inorganic pollutants such as heavy metals. Depending on the particle size that can be retained, different types of membrane filtration such as nanofiltration, ultrafiltration, and reverse osmosis can be utilized for heavy metal uptake from wastewater.
Ultrafiltration (UF) utilizes permeable membrane to separate heavy metals, macromolecules and suspended solids from inorganic solution on the basis of the pore size ranging from 5 to 20 nm and molecular weight of the isolating compounds (1000– 100,000 Da) [54]. Based on the membrane characteristics, UF can achieve more than 90% of removal efficiency with a metal concentration (10 - 112 mg/L) at pH ranging between 5 and 9.5 and at pressure (2–5 bar) . UF offers some advantages such as smaller space requirement and a lower driving force owing to its high packing density.
Polymer-supported ultrafiltration (PSU) process adds water soluble polymeric ligands to bind metal ions and form macromolecular complexes by generating a free targeted metal ions effluent [55]. PSU technology has the dvantages of the low-energy requirements involved in ultrafiltration, the reaction kinetics is very fast and higher separation selectivity of selective bonding agents in aqueous solution.
Another similar technique, complexation–ultrafiltration, confirms to be a promising alternative to technologies depend on ion exchange and precipitation. Using water-soluble metal-binding polymers in combination with ultrafiltration (UF) is a hybrid approach in order to concentrate selectively and to recover heavy metals in the solution. In the complexation – UF process cationic forms of heavy metals are first complexed by a macro-ligand in order that increasing their molecular weight with a size larger than the selected membrane pores [56, 57]. The advantages of complexation–filtration process iclude high selectivity of separation because of the use of a selective binding and low-energy requirements involved in these processes. Water-soluble polymeric ligands have shown to be powerful substances in order to separate trace metals from aqueous solutions and industrial wastewater through membrane processes.
Reverse osmosis (RO) is a separation process using pressure in which solution is forced through a membrane that keeps the solute on one side and allows the passage of the pure solvent to the other side. The membrane here is semi-permeable, meaning it allows the passage of solvent but not for metals. The reverse osmosis membranes have a dense barrier layer in the polymer matrix where most separation takes place. Reverse osmosis can remove many types of ions and molecules from solutions, including bacteria, and is used in both industrial processes. A diffusive mechanism is involved in reverse osmosis process, so that separation efficiency is dependent on pressure, concentration of the solute, , and water flux rate [58].
F. Electrodialysis:
Electrodialysis (ED) is a membrane separation uses electric potential to passe ionized species in solution through an ion exchange membrane. The membranes are plastic materials thin sheets with either cationic or anionic characteristics. When a solution containing ionic species passes through the compartments of the cell, the anions migrate toward the anode while the cations migrate toward the cathode, crossing the anion exchange and cation-exchange membranes [59]. A disadvantage of this process include membranes replacement and the corrosion process [60]. Using membranes with higher capacity of ion exchange resulted in better cell performance. Effects of temperature, flow rate,and voltage at different concentrations by using two types of commercial membranes, using a laboratory ED cell, on the removal of lead were studied [61]. The princible of Electrodialysis process is illustrated in Figure (5). Results show that increasing temperature and voltage improved cell performance and separation percentage decreased as the flow rate increasing. This provides advantages for the treatment of highly concentrated wastewater laden with heavy metals to recovery undesirable impurities from water.
Figure (5): Electrodialysis principles [62] CM – cation exchange membrane, D-dialute chamber, e1 and e2-electrode chambers, AM-anion exchange membrane and K-concentrate chamber
G. Adsorption:
Biosorption is another technique that can used for elimination of heavy metals from wastewater. Sorption process is defined as transfer of ions from solution phase to the solid phase, really describes a group of processes, which includes adsorption and precipitation reactions. Adsorption has become one of the alternative remidation techniques for wastewater. Basically, adsorption is a mass transfer process and substances bound by chemical and or physical interactions to solid surface [63-65]. All adsorption mechanisms are dependent on solid-liquid equilibrium and on mass transfer rates. A dsorption could be divided into the following types, depending on the types of intermolecular attractive forces [66, 67].
Physical adsorption:
It is a process in which binding of adsorbate on the surface of adsorbent as a result of Van der Waals forces of attraction or hydrogen bonding. Physical adsorption can only be occurred in the low temperature environment and under appropriate pH conditions.
Chemical adsorption:
A strong interaction arise from chemical reaction between the adsorbate and the adsorbent molecules is involved. This interaction produces new types of electronic bonds (Covalent and Ionic).
Mechanism of adsorption:
generally, the main steps involved in adsorption of contaminates on solid adsorbent are:
1.Transfer of the metal ion from the bulk of solution to the outer surface of the adsorbent.
2. Internal mass transfer by pore diffusion from outer surface of adsorbent to the inner surface of porous structure.
3. Adsorption of adsorbate onto the active sites of the adsorbent pores
4. The overall adsorption rate is determined by either intra particle diffusion or film formation or both as the last step of adsorption are very fast as compared to the other two steps.
The parameters which have been established for optimizing the use of adsorbent in wastewater treatment include [34]:
1. Nature of adsorbent and adsorbate.
2. Metal concentration.
3. pH and temperature of the aqueous solution.
4. Kinetics of adsorption.
5. Adsorption isotherm.
6. The time of contact.
Various low-cost adsorbents, derived from natural material, agricultural waste, industrial by-product, or modified biopolymers are found to be more promising and encouraging in heavy metal removal owing to various considerations as follow [36, 63].
(I)They are economical, (II) its metal selectivity, (III) they are regenerative, (IV) toxic production of sludge not present (V) metal recovery and (VI) its high effectiveness.
Using activated carbon in water and wastewater remidation has been directed towards organics removal [40] . Research efforts on removal of inorganics by activated carbon, specifically metallic ions, have been markedly limited [40] . selective adsorption by red mud [68], coal [69], photocatalyst beads [70], nano-particles[71], fertilizer industrial waste [72], biomass [73], activated sludge biomass [74], algae [75, 76] etc. has generated increasing excitement.
Industrial by-products such as fly ash [77] iron slags, waste iron [78] , hydrous titanium oxide [79, 80] ,can be chemically modified to enhance its removal performance for metal elimination from wastewater.
It was reported that [81, 82] for the disposal of heavy metals from industrial waste effluent has been focused on the use of agricultural by-products as adsorbents through biosorption process. New resources such as rice husk, coconut shell, pecan shells, rice straw, maize cob or husk, jackfruit, hazelnut shell, rice husk,…etc can be used as an adsorbent after chemical modification or conversion by heating into activated carbon or biochar for heavy metal uptake. They found that the maximum metal removal occurred by those biomass due to containing of cellulose, lignin, carbohydrate and silica in their adsorbent [83] .
Biopolymers are posse a number of different functional groups, such as amines and hydroxyls, which increase the efficiency of metal ion uptake [84] . They are widely use in industrially as they are able to lower the concentrations of transition metal ion to sub-part per billion concentrations. New polysaccharide-based-materials are described as biopolymer adsorbents (derived from chitosan, starch and chitin) for the elimination of heavy metals from the wastewater. The sorption mechanisms of polysaccharide-based-materials are complicated and depend on pH [84]. Also hydrogels, which are cross linked hydrophilic polymers, are widely used to purify wastewater. The removal is mainly governed by the water diffusion into the hydrogel, carrying the heavy metals inside especially in the nonexistence of strongly binding sites. Maximum binding capacity increases with higher pH because of polymerization/cross linking reaction.
1.2.2. Biological Methods:
Biological removal of heavy metals in wastewater involves the use of biological methods for the elimination of pollutants from wastewater. In this processes microorganisms play an important role of settling solids in the solution. Activated sludge, stabilization ponds, trickling filters are widely used for wastewater purification. Activated sludge is considered the most common option uses microorganisms in the treatment process to break down organic matter with agitation and aeration, and then allows solids to settle out. Bacteria-containing “activated sludge” is frequently re-circulated back to the aeration basin to increase organic decomposition rate. In biological systems, most of the research on heavy metals removal has been oriented towards the suspended growth activated sludge process. Trickling filters which consist beds of coarse media (often plastic or stones) 3-10 ft. deep help to grow microorganisms. Wastewater is sprayed into the air (aeration), then allowed to trickle through the media and microorganisms break down organic matters in the wastewater. The drain of trickling filters at the bottom and the wastewater is collected and then undergoes sedimentation. Lagoons or stabilization ponds are cheap, slow and relatively inefficient, biological method that can be used for different types of wastewater. They depend on the interaction of sunlight, microorganisms, algae, and oxygen [40].
1.3. Evaluation of heavy metals removal processes:
Although all the techniques of heavy metal wastewater treatment can be employed to eliminate heavy metals, they have their latent advantages and limitations. Table (2) indicates heavy metals uptake from aqueous solutions has been traditionally carried out by chemical precipitation because it is a simple process and cheap capital cost. However, chemical precipitation is ordinarily adapted for treating wastewater containing heavy metal ions with high concentration and it is ineffective with low metal ion concentration. Chemical precipitation is considered as not economical and can produce large sludge amount to be treated with great drawbacks [85].
Ion exchange has been commonly applied for the removal of heavy metal from wastewater. However, the resins of ion-exchange must be regenerated by chemical reagents when they are exhausted and the regeneration can cause serious secondary contamination. And it is expensive, particularly when treating a large amount of wastewater containing low concentration heavy metal, so that they cannot be used at large scale.
Adsorption is a common method for the uptake of heavy metals from low concentration aqueous solutions containing heavy metal. The activated carbon high cost limits its use in adsorption. Many varieties of adsorbents with low-cost have been developed and tested for heavy metal ions uptake. However, the efficiency of adsorption depends on the adsorbents type. Biosorptio of heavy metals from aqueous mediums is considered as new method that has proven very promising for the removal of heavy metal ions from wastewater [85].
The technology of membrane filtration can separate heavy metal ions with high efficiency, but its drawbacks such as high cost, low permeate flux process complexity, and membrane fouling have limited their use in heavy metal ions uptake.
Coagulation-flocculation technique can be employed for heavy metal wastewater remidation, the advantages of this method are dewatering and good sludge settling of the produced sludge. large sludge volume generation and chemical consumption are the limitations of this technique.
Table (2): The main advantages and disadvantages of the various physico-chemical methods for treatment of heavy metal in wastewater.
Treatment method
Target of removal
Advantages
Disadvantages
References
Chemical precipitation
Coagulation–flocculation
Dissolved air flotation
Ion exchange
Ultrafiltration
Nanofiltration
Reverse osmosis
Adsorption with new adsorbents
Heavy metals, divalent metals
Heavy metals and suspended solids.
Heavy metals and suspended solids
Dissolved compounds,
cations/anions
High molecular weight compounds (1000–10000 Da)
Hardness ions such as Ca(II) and Mg(II) and sulphate salts
Organic and inorganic compounds
Heavy metals
Low capital cost, simple operation
Shorter time to settle out suspended solids, improved sludge settling.
Low cost, shorter hydraulic retention time
No sludge production, less time consuming
Smaller space requirement
Lower pressure than RO (7–30 bar)
High rejection rate, able to withstand high temperature
Low-cost, easy operating conditions, having wide pH range, high metalbinding capacities
Sludge generation, extra operational cost for sludge disposal.
Sludge production, extra operational cost for sludge disposal.
Subsequent treatments are required to improve the removal efficiency of heavy metal
Not all ion exchange resin is suitable for metal removal, high capital cost
High operational cost, prone to membrane fouling
Costly, prone to membrane fouling
High energy consumption due to high pressure required (20–100bar), susceptible to membrane fouling
Low selectivity, production of waste products
[86-89]
[87]
[87, 90]
[91, 92]
[91, 93]
[87]
[87, 91]
[81, 94, 95]
. Flotation presents several advantages over the more conventional methods, such as high removal efficiency, high selectivity of metal ions, low detention periods, low operating cost production of more concentrated sludge and high overflow rates [96]. Operation costs, high initial capital cost and high maintenance are the disadvantages of this process.
Electrochemical heavy metal wastewater treatment technologies are considered as rapid and well-controlled that require fewer chemicals, offer good reduction yields and generate less sludge. On the other hand, electrochemical methods involving high cost electricity supply and high initial capital investment, this restricts the technique development.
Biological technologies by using different low materials were found be very effective techniques with higher uptake percentage. Although biological techniques are low cost and friendly methods for the environment they require large areas and proper operation and maintenance [40].
Even though all above techniques can be employed for the remediation of heavy metal wastewater, it is important to remarkable that the selection of the most suitable remidation methods depends on the initial concentration of the metal, wastewater component, plant flexibility and accuracy, capital investment, operational cost and environmental impact, ...etc [60, 85].
1.4. Enviromental pollution with iron metal and its removal:
Iron is considered the second metal among the most abundant metals on the earth crust. In the periodic table of elements, iron occupies the 26th elemental position. Iron is existing in many forms in water as shown in Figure (6) [97]. Biologically it is a most crucial element for survival and growth of almost all living organisms. As it is the cofactor for many vital enzymes and proteins. It is one of the vital constituents of organisms like algae and of enzymes such as catalase and cytochromes, in addition to oxygen transporting proteins, such as myoglobin and hemoglobin. Because of iron inter-conversion between ferrous (Fe2+) and ferric (Fe3+) ions, it is regarded as an attractive transition metal for various biological redox processes owing. The iron source in surface water is anthropogenic and is associated with mining activities. Sulphuric acid production and the discharge of ferrous (Fe2+) occurs because of iron pyrites (FeS2) oxidation that are common in coal seams. [98-100]. The following equations represent the simplified oxidation reaction for ferrous and ferric iron [99]:
2FeS2 + 7O2 + 2H2O 2FeSO4 + 2H2SO4 (ferrous)
4FeSO4 + O2 + 10H2O 4Fe(OH)3 + 4H2SO4 (ferric)
Mediated reactions of iron support the respiration process of most of the aerobic organisms. If it is not shielded correctly, it can catalyze the reactions involving radicals formation which can destroy biomolecules, tissues, cells and the whole organism. Iron poisoning has always been a subject of interest chiefly to pediatricians. Children are highly susceptible to iron toxicity as they are exposed to a maximum of products containing iron [101].
Figure (6): Classification of different forms of Iron presen in water.
Iron toxicosis occurs in four stages [100]:
The first stage which takes place after 6 hrs of iron overdose is noticeable by effects of gastrointestinal such as vomiting, diarrhea and gastro intestinal bleeding.
The second stage progresses within (6 - 24hrs) of overdose and it is regarded as the inherent period, a period of apparent medical recovery.
The third stage happens between 12 to 96 hrs after certain clinical symptoms onset. This stage is characterized by shocks, tachycardia, lethargy, hypotension, metabolic acidosis hepatic necrosis, and sometimes death.
The fourth stage take place in betwwen 2 to 6 weeks of iron overdose. This stage is marked by the gastrointestinal ulcerations formation and strictures development.
Iron uptake excess is a serious problem in meat eating and developed countries and it increases the cancer risk. Workers who are highly susceptible to asbestos that contains almost 30% of iron are at high risk of asbestosis, which is the second most important reason for lung cancer. It is said that asbestos associated cancer is related to free radicals. Loose intracellular iron can also promote DNA destruction. Iron can initiate cancer mainly by the DNA oxidation process.
Iron salts such as iron sulfate, iron sulfate heptahydrate and iron sulfate monohydrate are of low acute toxicity when exposure is through dermal, oral and inhalation routes and hence they have been placed in toxicity category 3. Moreover, the Food and Drug Administration considered that iron salts are safe and their toxic effects are very much negligible.
Free radicals formation is the outcome of the toxicity of iron. During pathological and normal cell processing, byproducts such as hydrogen peroxide and superoxide are produced, which are regareded to be free radicals. These free radicals are actually neutralized by enzymes such as catalase superoxide dismutase, and glutathione peroxidase but the superoxide molecule has the capability to release iron from ferritin and that free iron reacts with more and more of hydrogen peroxide and superoxide forming free radicals with high toxicity such as hydroxyl radical. Hydroxyl radicals are dangerous as they can initiate lipid peroxidation, inactivate certain enzymes, cause DNA strand breaks and can depolymerize polysaccharides. This can occasionally lead to cell death.[100, 102, 103]
Tahir and Rauf [104] studied the Removal of Fe(II) from the galvanized pipe manufacturing industry wastewater by adsorption onto bentonite clay. The adsorption of Fe(II) from aqueous solutions over a concentration range from 80 to 200 mg/l, shaking time of 1–60 min, adsorbent dosage 0.02 – 2 g and pH of 3. The process of removal follows both the Langmuir and Freundlich isotherm models and also obey the first-order kinetics. The maximum removal (> 98%) was observed at pH of 3, 0.5 g of bentonite with initial concentration of 100 mg/l. The Fe(II) removal efficiency was also tested using wastewater from a galvanized pipe manufacturing industry. Higher than 90% of Fe(II) can be effectively removed from the wastewater by using 2.0 g of the bentonite. The effect of cations (i.e. manganese, cadmium, lead, chromium, zinc, copper, nickel and cobalt) on the removal of Fe(II) was studied in the concentration range of 10–500 mg/l. All the added cations reduced the Fe(II) adsorption at high concentrations except Zn. Column studies have also been investigated using a certain concentration of wastewater. More than 99% recovery has been attained by using 5 g of the bentonite with nitric acid solution (3 M).
Das et al. [105] in a study investigated that the traditional method of using ash for disposal of iron from groundwater can eliminate iron to desired level without increasing the pH behind the acceptable limit. The banana pseudostem ash is among the different plant ashes used for iron removal. It has been found to be most appropriate for iron elemination. The ash improves iron uptake. The designed iron elemination system is expected to be convenient for household use. The optimum values of the different parameters for iron removal are 200–300 mg L−1 ash, 1.0 L h−1 rate of filtration and time of residence (1h) for groundwater having 2.20 ppm iron concentration. For groundwater having higher [Fe], the amount of ash can be increased and can be decreased gradually throughout continuous use. The technique has the advantages of low manufacturing cost, almost nil recurring cost, viz., simplicity in use, no electricity requirement, and increasing the essential minerals such as K, Ca in the treated water.
Al-Anber et al. [106] in their study the batch removal of Fe3+ from aqueous model solution under different experimental conditions using Jordanian natural zeolite (JNZ) has been investigated. The contact time influences, initial concentration of metal, temperature and concentration of adsorbent dosage have been studied. The adsorption efficiencies are found to be residence time dependent, increasing the contact time in the range between 1 and 150 min. The sorption equilibrium has achieved between 60 and 150 min. The optimum adsorption has occur at 30°C of temperature. The equilibrium adsorption capacity of JNZ adsorbent used for Fe3+ were evaluated and extrapolated using Freundlich and Langmuir isotherm models and the experimental data are found to fit Langmuir isotherm more than Freundlich isotherm.
Ghosh and his team [107] reported the results of astudy on electrocoagulation (with electrodes from aluminum) for iron Fe(II) elemination from aqueous medium. The removal of Fe(II) was composed of two principal steps; (a) oxidation of Fe(II) to Fe(III) and (b) subsequent Fe(III) disposal by the freshly formed aluminum hydroxides complexes by adsorption/surface complexation followed by precipitation. Experiments were executed with various current densities ranging between 0.01 and 0.04 A/m2. Other parameters such as salt concentration, pH and conductivity were maintaned constant as per tap water quality. It was observed that as the current densities increase, the elemination of Fe(II) increased. Satisfactory iron removal of around 99.2% was attained at the end of 35 min of operation from 25 ppm initial Fe(II) concentration.
Vasudevan et al [108] reported the results of astudy on the uptake of Iron from drinking water by electrocoagulation using galvanized iron as the cathode and magnesium as the anode. Experiments were done as a function of current density, pH and temperature. By using both the Langmuir and the Freundlich isotherm models, the adsorption capacity was estimated. The results demonstrated that the maximum efficiency of removal of 98.4% was obtained at 0.06 A dm– 2 of current density and pH of 6.0. The adsorption of iron was better illustrated by fitting the Langmuir adsorption isotherm, which suggests adsorbed molecules monolayer coverage. The adsorption process followed a kinetics model of second-order. Temperature studies illustrate that adsorption was endothermic and spontaneous in nature.
Bulai and Cioanca [109] in astudy found that Purolite S930 is an effective sorbent for Iron (II) ions uptake from aqueous model solutions in different operating conditions. The Iron (II) elemination percent has a maximum at pH 5.0 and increases with contact time and resin dose increasing and decreases with solution initial concentration increasing.
Wang [110] investigated on crushed concrete and limestone removed Fe(II) from synthetic groundwater in laboratory columns. The results shown that achieving average Fe(II) uptake of greater than 216 (approximately 133 L treated) and 99% over 288 (approximately 172 L treated) pore volumes, for crushed concrete and limestone, respectively. Calcium siderite which formed in limestone columns as a form of precipitate; this formation had no significant effect on porosity of the system, but may have impeded Fe(II) remidation by limiting available surface area for adsorption. Results suggested that field-scale passive iron removal systems, using similar materials,merit exploration. Because differences are expected, pilot-scalefield tests are warranted.
Nandeshwar et al. [111] in astudy foucsed on iron uptake from real wastewater samples of Nag River, India using Green activated carbons from different waste materials such as orange peels, sawdust, C. procera leaves and coconut shells. All the selected waste materials were carbonized in muffle furnace and activated using various agents such as HCl, HNO3, and H2SO4. The results showed that all adsorbents have the potential capacity to separate iron, which further highly increases after its activation. The most promising green adsorbents were found to be orange peels and HCl was the best activating agent. The order of iron uptake from wastewater is: orange peels then coconut shells then sawdust then C. procera leaves. Similarly it was found that charcoal activated with HCl can separate around 77–90% iron followed by HNO3 (70–80%) and H2SO4 (58–75%).
Vries et al. [112] conducted astudy on iron separation from water by using rapid sand filtration. A model has been developed that takes into account the main properites of (submerged) rapid filtration: the water quality parameters of the influent water, marked pH, concentration of iron(II), homogeneous oxidation in the supernatant layer, surface sorption and heterogeneous oxidation kinetics in the filter, and adsorption characteristics of the filter media. Adsorption isotherm data collected from different Dutch remidation places show that Fe(II) adsorption may vary strongly between them, but generally increases with higher pH. The model has a sensitivity for (experimentally) determination of adsorption parameters and the heterogeneous oxidation rate.
Wang et al. [113] in astudy focused on Effects of solution chemistry on the removal reaction between Fe(II) and calcium carbonate-based materials by using a permeable reactive barrier consist of calcium carbonate-based materials (CCBMs), such as limestone. There is no significant effect on the uptake of Fe(II) by limestone from pH 7 to 9. Na+ significantly affected elemination of Fe(II) at levels of 100 mg/L and above. Ca2+ and Mn2+ showed effect on removal as low as 10 ppm Ca2+ and 5 ppm Mn2+. natural organic matter (NOM) premixed with Fe(II) (10 ppm Dissolved organic carbon (DOC) ) resulted in final Fe(II) levels above GCTL (groundwater cleanup target level). NOM retained 0.05 mg Fe(II)/mg for 2/3 sources and 0.032 mg/mg for 1/3.
Indah and Helard [114] conducted astudy on Evaluation of Iron and Manganese-coated Pumice from Sungai Pasak, West Sumatera, Indonesia for Fe (II) and Mn (II) Removal from aqueous model solutions. The effect of soaking time for iron and manganese coating was studied and as comparison. The experiments were performed in batch mode at room temperature between 20 and 25 oC, pH 7; adsorbent dose of 10 g/L; adsorbent diameters of 0.30-0.50 mm; 90 minutes of soaking time and100 rpm of agitation speed. The results showed that the optimum soaking time for manganese coating and iron for removal of Fe (II) and Mn (II) was 100 hours. Iron-coated pumice showed to have high removal efficiency compared to uncoated and manganese-coated pumice. More than 84% of Fe(II) with 15 ppm initial concentration was removed by 10 g/L iron-coated pumice, while by using uncoated and manganese-coated pumice, the elemination efficiencies were less than 75% . The desorption study noticed that up to 20% of Fe (II) was recovered from the three kinds of pumice adsorbent. Overall research indicated that pumice from Sungai Pasak may be a promising adsorbent for iron disposal from water and wastewater.
1.5 Enviromental pollution with lead metal and its removal:
Lead is a highly toxic metal whose widespread use has give rise to comprehensive environmental pollution and problems of health in many world parts. Lead is a bright silvery metal, slightly bluish in a dry atmosphere. It begins to tarnish with air contact, thereby forming a complex mixture of compounds, depending on the given conditions. Figure (7) shows various sources of lead pollution in the environment [115].
The lead exposure sources involve mainly industrial processes, smoking and food, drinking water and domestic water sources. The lead sources were gasoline and house paint, which has been extended to plumbing pipes, lead bullets, storage batteries, pewter pitchers, faucets and toys [116]. larger than 100 to 200,000 tons of lead per year is being emitted from car exhausts in the US. Some is taken up by plants, fixation to soil and flow into water bodies, hence human exposure of lead in the general population is either owing to drinking water or food. Lead is an extremely toxic heavy metal which disturbs different physiological processes of plant and unlike other metals, such as, copper, manganese and zinc, it does not play any biological functions. A plant with high concentration of lead fastens the reactive oxygen species (ROS) production, resulting in damage of lipid membrane that finally leads to destruction of photosynthetic and chlorophyll processes and suppresses the plant overall growth [117]. Some research stated that lead is capable of suppressing the tea plant growth by reducing biomass and debases the tea quality by changing the quality of its components [118]. Even at low concentrations, lead remediation was found to cause large instability in ion uptake by plants, which in turn leads to significant metabolic changes in photosynthetic capacity and ultimately in a strong inhibition of plant growth.
Figure (7): Various sources of lead pollution in the environment.
Poisoning of Lead was considered to be a classic disease and the marks that were seen in children and adults were fundamentally attached to the gastrointestinal tract and the central nervous system [119]. Lead poisoning can also take place from drinking water. The pipes which carry the water may be made of lead and its compounds which can pollute the water [120]. According to the Environmental Protection Agency (EPA), lead is regarded a carcinogen. Lead has large effects on different parts of the body. Distribution of Lead in the body initially based on the blood flow into various tissues and almost 95% of lead is precipitated in the form of insoluble phosphate in skeletal bones [121]. Toxicity of lead, also called lead poisoning, can be either chronic or acute. Acute exposure can result in headache, abdominal pain, appetite loss, renal dysfunction, vertigo, sleeplessness, arthritis, hallucinations, hypertension and fatigue. Acute exposure chiefly occurs in the work place and in some manufacturing industries which make use of lead. chronic exposure of lead can cause psychosis, autism, weight loss, allergies, mental retardation, dyslexia, hyperactivity, kidney damage, paralysis, muscular weakness, birth defects, brain damage, and may even lead to death [122].
Eventhough lead toxicity is preventable it still remains a dangerous disease which has effect on most of the organs. The plasma membrane moves into the brain interstitial spaces when the blood brain barrier is exposed to great levels of lead concentration, leading a condition called edema. It disrupts the intracellular second messenger systems and alters the the central nervous system functioning, whose protection is highly important. Domestic and environmental lead ions sources are the main reason of the disease but with appropriate precautionary measures it is possible to reduce the risk correlated with lead toxicity [120].
Generally, Impact of lead exposure in humans has been known to cause wide variety of health problems such as [123] :
• Various forms of blood disorders and Anemia
• Rapid deterioration of brain and the nervous system
• fertility decreasing both in men and women
• Failure of the kidney
• Alzheimer disease
Many studies have been reported for lead elemination from aqueous solutions. Pala and Dursun [124] studied that the results of a study on adsorption of Pb (II) ions from artificial contaminated tap water by using a natural zeolite (Clinoptilolite). Clinoptilolite mineral which has mesh size of 25-140 was used by activating with HCl, and the efficiencies of lead ion disposal were evaluated. Experiments were occured under laboratory batch conditions were run at different values of pH, temperatures. The highest efficiency of removal was found as about 87% at pH 5. In similar way, experiments were done at different temperature values, and the utmos efficiency was achieved at 30oC. The efficiency obtained under these conditions was 89.95%. The highest lead disposal efficiency was achieved with shaking speed of 200 rpm.
Mavropoulos et al. [125] in their study found that the composite of hydroxyapatite-alginate was effective in the elemination of lead ions and lead phosphate nanoparticles from high-polluted simulated gastric fluid. The cross-linked polymer chain had a double role: (i) keep Pb2+ ions and lead phosphate nanoparticles bounded to the surface of bead, impeding their bioavailability in stomach fluid; and (ii) delay dissolution of HA in the stomach acidic conditions, confirming that an excess of Ca2+ will not be released to simulated gastric fluid. Desorption studies in simulated enteric fluid stated that lead stayed immobilized in the calcium phosphate phase in the intestinal tract. These results indicate HA–alginate composite as effective system for heavy metals disposal from polluted gastric and enteric human fluids, reducing its adsorption by the human body.
Meski and his team [126] showed in their work with hydroxyapatite prepared from the egg Shell , that the carbonate hydroxyapatite prepared from egg shell (CHAPF) represents the highest capacity for Pb2+ ions adsorption from aqueous solution. It has been found that the initial adsorption rate was high. The sorption process obey the model of Langmuir isotherm with low temperature dependency and high adsorption capacities. The thermodynamic functions were calculated, and it can be concluded that the Pb2+ adsorption over CHAPF is an exothermic and spontaneous process. The adsorption was greatly pH dependent, with a high uptake of lead at pH = 3. These results show that the lead uptake by CHAPF was very sensitive to the initial concentration of Pb2+ in aqueous solution. For the high concentrations [(500 to 700) mg·L-1], two stages were observed: Pb2+ ions adsorption on the CHAPF surface and an ion exchange reaction between Ca2+ of CHAPF and Pb2+ ions in aqueous model solution.
Shrestha et al. [127] conducted a study on lead (II) disposal from aqueous solutions using prepared activated carbon. Two series of carbon have been synthesized from Lapsi seed stones by treating with concentrated H2SO4 and HNO3 in a mixture with H2SO4 in the ratio of 1:1 by weight for disposal of metal ions. pH 5 was the optimum pH for lead adsorption. For the equilibrium isotherms description, the adsorption data were better fitted with the Langmuir adsorption equations than Freundlich equation. The maximum adsorption capacity of Pb (II) on the produced activated carbons was 277.8 mg/g with a mixture of HNO3 and H2SO4 and 423.7 mg/g with H2SO4. The waste material used in activated carbons preparation is readily available and cheap. Therefore the carbons synthesized from Lapsi seed stones can work as potential low cost adsorbents for the elemination of Pb (II) from water.
Jalali [128] investigated astudy on stalk of Sunflower, an agricultural waste, acts as an adsorbent for the cadmium and lead disposal from aqueous solutions. Adsorbent was synthesized by washing residue of sunflower with deionized water until the solution become colorless. The results stated that the adsorbent has good sorption potential and maximum removal of metal was detected at pH 5. Within 150 min of operation about 97 of Pb ions were eleminated from the effluents. Curves of lead sorption were well fitted to the modified two-site Langmuir model. Lead adsorption capacities at optimum operation conditions were 182 mg/g. The kinetics of Pb ions adsorption from aqueous model solutions were also analyzed. The experimental data were foud to be fitted to pseudo-second-order kinetic model. fitted models (R2 > 0.999). The maximum adsorption capacity for Pb(II) ions adsorbed onto entrapped silica nanopowders was evaluated to be 83.33 mg/g.
Soltani and his team[129] conducted astudy on adsorption of Pb(II) ions from the aqueous solution by using entrapped silica nanopowders within calcium alginate in order that determination the thermodynamic, isotherm and kinetic of the adsorption process. According to the results, an initial pH of 5.0 was found to be optimal for the Pb(II) ions adsorption. The capacity of adsorption reached to 36.51 mg/g with increasing the contact time to 180 min at 50 ppm as initial Pb(II) ions concentration. However, the equilibrium contact was estimated to be 90 min owing to no significant increase in adsorption effeciency after this time. The results of studies stated that the isotherm of Langmuir and pseudo-second order model of kinetic were the best.
Yarkandi (2014) [130] carried out batch experiments for lead separation from waste water using natural american bentonite and activated carbon. The results show that the amount of Pb++ adsorption increases with solution pH, initial concentration of metal ion and contact time but decreases with temperatures and amount of adsorbent. The adsorption process has well fit pseudo-second order kinetic model. Langmuir and Freundich adsorption isotherm models were found to be applicable to the adsorption process where both were applies to analyze adsorption data. Thermodynamic parameters e.g. ΔH°, ΔG° and ΔS° of the adsorption process was found to be endothermic. Finally it can be seen that activated carbon was found to be less effective for disposal of Pb+2 ions than bentonite.
Bartczak et al [131] in a study investigating the lead (II) ions adsorption from aqueous model solution on peat as adsorbent with low-cost, observed that The sorption capacities of peat with respect to lead(II) ions was 82.31 mg(Pb2+)/g. Slightly well results were achieved with adsorption efficiency reached 100% just after 3 min (15 and 30 ppm), 5 min (50 ppm) or 15 min (100 ppm) of the process. It indicates the utmost affinity of the peat surface for lead ions. the optimum adsorbent mass was found to be 5 g/L. To prevent metal precipitation as hydroxides and including the obtained results pH = 5 confirmed as an ideal.
Sangeetha et al [132] investigated a study on lead ions disposal from aqueous solution by using novel hydroxyapatite/alginate/gelatin composites. Pb2+ elemination ability of wet precipitation synthesized biosorbents HA/Alg and HA/Alg/Gel has been investigated for different dosages. Complete disposal was obtained from 7 to 24 hours by both the adsorbents. The sorption kinetics were found to be best fit to the pseudo-second order equation and the equilibrium well followed Langmuir isotherm model. The elemination capacity was higher for lower dosage studied and the rate of disposal was higher for the higher dosage studied.The sorption mechanism involved was dissolution/precipitation, ion exchange and surface complexation process. according to The results, it can be seen that both the composites under study are potential candidates for Pb2+ disposal and precisely gelatin enhanced the maximum sorption than the alginate alone which is composed with the hydroxyapatite.
Cheraghi et al. [133] showed in their work with waste tea leaves, upon parameters optimization like initial metal concentration, Temperature, pH and adsorbent amount, that maximum uptake efficiency was achieved at pH 6. Also, as the initial metal concentration decreased the adsorption of Pb (II) ions increased. the equilibrium adsorption isotherm data fits well with the Langmuir isotherm model and its calculated maximum adsorption capacity of monolayer was 166.6 mg/g at 25±0.1˚C. The sorption kinetics were found to be best fit to the pseudo-second order equation.
Kanyal and Bhatt [134] reported the results of astudy on adsorption of Pb (II) ions from waste water by using household waste as an adsorbent. Banana peels, Pumpkins and Chicken eggshells are considered as good adsorbents for elemenation of heavy metals from polluted water. The effects of various parameters such as agitation speed, pH and residence time were examined and best results were observed at pH 7, 90 mins and 100 rpm. The results stated that household waste usage such as these can be act as a good biosorbent for disposal of heavy metals on a large scale and establish effective, and inexpensive techniques in wastewater remidation.
Wang et al. [135] carried out a study for selective disposal of lead and other metals such as cadmium and copper from wastewater by gelation with alginate for effective recovery of metal. The results evaluated that gels can be formed speedily between the metals and alginate in lower than 10 min and the rates of gelation fit well with the pseudo second-order kinetic model. The optimum ratio of dosing of alginate to the metal ions was found to be between 2:1 and 3:1 for Pb2+ elemination and around 4:1 for Cu2+ and Cd2+ elemination from wastewater, and the metal disposal efficiency by gelation increased with the solution pH. Alginate has a higher affinity of gelation toward Pb2+ than Cd2+ and Cu2+, which permitted a selective uptake of Pb2+ from the wastewater in the existing of Cd2+ and Cu2+ ions.
Liu et al. [136] focused on the disposal of lead ion from waste water using hydroxyapatite scaffolds synthesized from scales of fish. Powder of fish scale obtained from Tilapia fish (Oreochromis mossambicus) was used for preparing scaffolds for lead removal. The maximum adsorption capacities (qmax) were 344.8 mg/g and 208.3 mg/g in solutions pH of 2.2 and 5, respectively. More than 99.9% of the lead ion was eleminated after 20 min.
Al Lafi et al. [137] conducted astudy on Lead removal from aqueous aolutions by polyethylene waste/nano-manganese dioxide composite . The adsorption results investigated that the synthesized adsorbent can effectively eleminate Pb+2 ions from aqueous solutions, with a maximum adsorption capacity of 50.5 mg/g. Almost 60 % of the initial Pb+2 concentration were adsorbed within the first hour, and it was concluded that 2 h was the optimum time for Pb+2 elemination . A pH value of 5.0 was determined as an optimum and was used for the rest of this study. Regeneration of the composite can be performed using 0.5 mol/L HCl solution with Pb+2 percentage of recovery reached 95 %. It also efficiently adsorbed Pb2+ after five sorption/desorption cycles with 84 % as percentage of removal.
Heraldy [138] conducted a study of sorption of Pb(II) from the aqueous mediums using biosorbent synthesized from waste of tomato and residue of apple juice (AR). The optimum conditions for maximum removal percentage of Pb(II) by biosorbents were found to be 0.1 g of sorbent at pH 4.0 and 90 min contact time for tomato waste and 60 min for AR. The experimental data were found to be well matched with Freundlich than Langmuir isotherm model. A kinetic study showed that Pb(II) sorption follows the pseudo-second-order kinetics, which confirms that AR and waste of tomato biosorbents are 108 and 152 mg/g, respectively.
1.6 Role of alginate in heavy metals removal:
Alginate is belong to the anionic polymers family which is naturally occurring usually produced from brown seaweed, and has been widely investigated and used for many biomedical applications, due to its relatively low cost, low toxicity, biocompatibility, and mild gelation by divalent cations addition such as Ca2+ [139]. Alginates are a polysaccharides composed of variable ratios of β-D-mannuronate (M) and its C-5 epimer α-Lguluronate (G) linked by 1–4 glycosidic bonds (Fig.8). In the 1880s, alginates were first separated from brown seaweeds, and its production for commerce started in the early 20th century.
The production of alginate can be carried out by two bacteria genera, Azotobacter and Pseudomonas and various genera of brown seaweed and [140]. Alginate which is available in commerce can usually extracted from brown algae (Phaeophyceae), including Laminaria digitata, Macrocystis pyrifera, Laminaria japonica, Ascophyllum nodosum and Laminaria hyperborea by remidation with aqueous alkali solutions, usually with NaOH. The extract is filtered, and in order to precipitate alginate either calcium or sodium chloride is added to the filtrate. By treatment with dilute HCl, the alginate salt can be converted to alginic acid. After further purification and conversion, power of water-soluble sodium alginate is generated. [139, 141]
Figure (8): Chemical structure of alginate. M – mannuronate residues, G – guluronate residues.
The synthesised alginates were found to be extremely effective in uptake of heavy metal ions from aqueous model solutions. Removal of more than 93% Cr(VI) ions was obtained from aqueous solution in batch process using this type of biosorbent. [142, 143]
Alginate is regarded a biopolymer with many applications in food industry, drug delivery systems, cosmetics and cell encapsulation. In wastewater remediation could play a significant role in disposal heavy metal ions due its advantages, such as biodegradability, facile obtaining procedure, economical, biocompatibility and environmental friendly [144].
NiŃa et al. [144] in astudy found that calcium alginate microparticles has a good affinity for the divalent cations. The heavy metals adsorption was examined as a function of residence time between the samples of synthetic wastewater and the alginate and the polymeric beads morphology. The microparticles of calcium alginate were synthesized using a laboratory procedure, sodium alginate aqueous solution was dropped in a solution of calcium chloride. The polymeric microparticles which have a controlled porosity seem to be an appropiate alternative to develop a aprocess for elemination of heavy metal from industrial wastewater.
Singh et al. [145] investigated astudy on effective removal of Cu2+ ions from aqueous medium using alginate as biosorbent. Maximum removal of Cu2+ ions (85.3%) from aqueous medium was observed at pH 5.5, alginate dosage of 2.5% and initial copper concentration of 275 ppm with 50 min as agitation time. Thus, the reultant experimental data has been fitted well with both Langmuir and Freundlich isotherm models.
1.7. Heavy metals removal by bentonite and kaolin clays:
Clay is one of potential good adsorbent substitutes to other adsorbent owing to its layered structure, large surface area, mechanical and chemical stability and high capacity of cation exchange. The existence of two acidity types, Lewis and bronsted in clays increases the clays adsorption capacity. Aluminum oxides and clay minerals such as bentonite and kaolin, are the most wide-spread minerals of the earth crust which are known to be good adsorbent of different metal ions, organic ligands and inorganic anions [146, 147].
1.7.1. Bentonite:
Bentonite is considered a strong candidate as an adsorbent for heavy metal elimination due to its abundance and its low cost. Bentonite as a representative clay mineral is a clay chiefly consist of montmorillonite, a 2:1 type of aluminosilicate Bentonites are extremely valued for their sorption [146, 147].
Zia and his team [148], evaluated the removal of Pb2+ and Cd2+ by using Methionine modified bentonite/Alginate (Meth-bent/Alg) nano composite. The desorption study presents that 99% of the adsorbed Pb2+ and Cd2+ can be desorbed by using oxalic acid (0.1 M) as eluting agent with regeneration ability up to fifth cycle effectively.
Wu et al. [149] reported the adsorption of Th4+ from aqueous solution by using Novel magnetic organo-bentonite-Fe3O4 (polysodium acrylate) (OB-Fe3O4 PSA) superabsorbent nanocomposites. OB-Fe3O4 PSA super absorbent could be regenerated through the desorption of Th4+ using HCl solution (0.1 mol/L) and the adsorption capacity was still greater than 3.6 mmol/g after five successive adsorption–desorption processes.
Tan and Ting [150] reported that plain alginate and alginate immobilized bentonite beads have good reusability potential. remidation with HCl (10 mM) successfully eluted 93.05% and 94.33% of the Cu2+ ions loaded onto plain alginate and alginate immobilized bentonite, respectively, after three cycles of sorption–desorption test. There was no significant difference in the percentage of Cu2+ desorped in the three sorption–desorption cycles for both the plain alginate (93.15%, 93.54% and 92.48%) and immobilized-bentonite (92.38%, 96.03% and 94.75%). This established high reusability of the developed immobilized bentonite without remarkable losses in their Cu2+ disposal capacities.
1.7.2. Kaolin:
Kaolin is considered a type of clay rock, which incloses some chemical elements such as Na, Al, Mg, Ti, Ca, Fe, Ka, Si , and so on. The silicon mass proportion is more than 50%. Kaolin can be divided into two classes, which is coal-series kaolin and nature kaolinite The monolayer crystal structure of kaolin is comprised of siliconoxygen tetrahedral sheets and aluminum-oxygen octahedral sheets, just like the molecular structure model of kaolin shown in Figure (9) [151]. Metal ions removal using kaolinite clay is depend on mechanisms of ion exchange and adsorption and kaolinite has a relative low capacity of cation-exchange (CEC) [3–15 meq/100 g of clay] and smaller surface area ranged between 10 and 20 m2/g [152].
Figure (9): The molecular structure model of the kaolin.
Li et al. [153] reported that a novel environmental friendly material, calcium alginate immobilized kaolin (kaolin/CA), which synthesized using a sol-gel method, have good effeciency for copper uptake from waste water. the experimental adsorption was described using the Langmuir isotherm, the maximum capacity of Cu2+ adsorption by the kaolin/CA reached up to 53.63 mg/g. The thermodynamic studies indicated that the adsorption reaction was found to be an endothermic and spontaneous process.
Yavuz and his team [154] investigated the elemination of heavy metals such as Cu(II), Ni(II), Co(II) and Mn(II) from aqueous solution using raw kaolinite. The sorption of these metals on kaolinite conformed to Langmuir adsorption equation. Langmuir Cm constants for each metal were found as 0.919 mg/g (Co), 10.787 mg/g (Cu), 1.669 mg/g (Ni), 0.446 mg/g (Mn), at 25 oC, respectively. Also, kinetic and thermodynamic parameters like entropy (ΔS), enthalpy (ΔH) and free energy (ΔG) were evaluted and indicated that heavy metal adsorption on kaolinite was an endothermic process and the process of adsorption was preferable at elevated temperatures.
Larakeb (2017) [155] evaluted the Zinc Removal from Water by Adsorpion on Kaolin and Bentonite clays. The kinetics of adsorption results showed that zinc disposal is max. with and 45.48℅ efficiency for kaolin after 60 min of residence time and after 20 min with 89.8 ℅ efficiency for bentonite. Adsorbent dose increasing from 0.5 to 8 g/l enhance zinc elemination efficiency for 5 ppm like an initial concentration. Zinc disposal efficiency by the two adsorbent decreases with rising of the initial Zn concentration from 2 to 20 ppm. pH of treatment has considerable effects on the retention rate of zinc. The efficiencies of Zn removal are noticeable at basic pH. Whatsoever reaction parameter tested, it appears that kaolin is less effective than bentonite.
1.8. Removal of heavy metals by gelatin:
It has been reported that gelatin has agood affinity for heav metals removal, individaly or combined in composites with other materials. Itabashi et al. [156] found that The good effect of copper elimination by gelatin was achieved by the foam treatment of this solution. And also lead was successfully eliminated by the same treatment. gelatin powder was used for the adsorption treatment to raise this effect before the foam treatment. About 99 % of copper in the range of pH between 6.5 and 7.3 and approximately 100% of lead at pH 7.0 was eliminated respectively. Hayeeye et al [157] in their study found that, The maximum capacity of adsorption of gelatin/ activated carbon for Pb2+ ions was obtained to be 370.37 mg g-1. The separation process for Pb2+ ions was found to be relatively rapid with 92.15% of the adsorption finished in about 5 min as residence time in batch conditions. Adsorption was achieved at pH value as low as 2.0 and maximum adsorption was observed at a pH of almost 5.
1.9. Hydroxy apatite from bio waste materials:
Hydroxyapatite (HAP, Ca10(PO4)6(OH)2), a naturally available form of calcium phosphate and a component of hard tissues, has been reported to work as an efficient ion uptake material for different heavy metals from aqueous medium owing to its low solubility of water and excellent reactivity. The high stability of HAP structure, along with its flexibility permit a high variety of exchanges (particularly Ca ions with divalent heavy metal ions, such as As, Cd, Cu, Zn, Pb, Co, Ni, Sb,U, Hg, of huge importance in the environmental science field [158-160]. HAP can be extracted from different biowaste such as, eggshells and bovine bones. It can be synthesized through various methods which can be generally divided into two major routes: solid state reaction and wet methods [161]. including sol-gel technique, wet precipitation, hydrothermal process, mechanochemical method. Depending on the used techniques, HAP with several morphologies, composition, specific surface and crystalline degree have been obtained and appear to have different effects on the mechanical properties, bioactivity and dissolution behavior in biological environment [160, 162].
1.9.1. Preparation of hydroxyl apatite by sol-gel technique:
Sol-gel method is used for obtaining HAP powder of fine particle (nanoparticle size). In the sol-gel method of the HAP the calcium compounds and phosphorus precursors are transformed through condensation and hydrolysis reactions to the amorphous gels, which are further converted to ceramics when heated at comparatively low temperature. The polycondensation and hydrolysis are not separated in time, but occurs simultaneously [163]. Ceramic materials prepared by sol-gel route present many advantages over the others, such as homogeneous composition, low synthesis temperature and high product purity. Additional advantage of the sol-gel method is its applicability for surface coating.
There is no any form of secondary environmental damage as a result of high biocompatibility and its slightly- alkaline pH. The efficiency of HAP in eliminating heavy metal ions extremely depends on ion nature, diameter, charge and concentration, in addition to the treated water properties (temperature, pH) [160]
Putra and his team [164] showed in their work with eggshell, that for batch adsorption studies , at 90 min equilibrium time, 0.1 g biomass dosage and pH 6 were optimum biosorption conditions for Zinc and Copper ions elimination from aqueous mediums.
Agarwal and Gupta [165] in their study with eggshells focused mainly on evaluting varying concentrations (5, 10, 20, 40, 100 mg/L) of lead and copper ; this study reported a 92% - 100% removal of Cu when 0.5 to 1.5 g of eggshells (adsorbent) was used against 5 and 10 ppm of Cu; and adsorption efficiency of 80% to 100% for Pb at the same concentrations.
Deydier et al. [166] conducted a study for elemination of Pb from effluents using alow-cost material from meat and bone meal combustion residue. This residue was regarded as an apatite-rich material and was used as a low-cost substitute of hydroxyapatite in lead elemination from water . the mechanism was found to be as in pure apatite: surface complexation and dissolution of calcium hydroxyapatite, followed by lead hydroxyapatite precipitation.
Rohaizar et al. [167] in astudy, observed that pH = 7 and 350 rpm as an optimum agitation rate were ideal for copper elemination from water.
Avram et al. [160] in a study investigated that low crystallinity HAP which prepared by the direct reaction of diammonium hydrogen phosphate and calcium nitrate at alkaline pH, can be successfully used in heavy metal disposal from mine wastewater. For all the 10 metals studied (Zn, Cd, Pb, Mn, Co, Fe,Cr, Cu, Ni and Al), their content was fastly reduced by contact with HAP under the legal allowable limits for wastewater discharge in natural environment. The ion exchange importance in sorption processes was revealed and the pseudo-2nd order kinetics of manganese ions sorption on HAP was estimated.
1.10. Heavy metals removal by silica nanoparticles (Diatomeous):
Silica is used widely in nanoparticles coatings which used in water purification techniques. Silica coating activates the NPs surfaces having various functional groups owing to the abundant existing of silanol groups on the silica layer. It also prevents leaching low pH situations of NPs. It also facilitates the NPs with non- specific moieties, highly and group specific ligands. Polymer layered silicate nanocomposites possess improved properties at low filler contents. At neutral pH, as the particle size increase, the acidity of Si NPs will increase resulting in 5 to 20% ionisation of silanol groups, causing attraction between anionic Si surface and cations by ion pairing [168].
Surface-functionalized nanoporous silica, often referred to as self-assembled monolayers on mesoporous supports (SAMMS), has previously presented the ability to act as very effective sorbents for heavy metal uptake in a range of environmental and aquatic systems [169]. Diatomite is belong to the siliceous rock family, silicon dioxide is the main constituent of it with the proportions up to 90%. Diatomeous has some advantages such as wear resistance, heat resistance non-toxic and large specific surface area,..etc. Diatomite is a sort of polyporous material. The diatomite porosity is up to 90%, which means that diatomite has great adsorbability [151].
Soltani et al. [129] conducted astudy on adsorption of Pb(II) ions from the aqueous solution by using entrapped silica nanopowders within calcium alginate in order that determination the thermodynamic, isotherm and kinetic of the adsorption process. According to the results, an initial pH of 5.0 was found to be optimal for the Pb(II) ions adsorption. The capacity of adsorption reached to 36.51 mg/g with increasing the contact time to 180 min at 50 ppm as initial Pb(II) ions concentration. However, the equilibrium contact was estimated to be 90 min owing to no significant increase in adsorption effeciency after this time. The results of studies stated that the isotherm of Langmuir and pseudo-second order model of kinetic were the best.
Karnib and his team [170] used a composite from Activated Carbon, Silica and Silica Activated Carbon. Silica/AC (2:3) composite showed the greatest elemination percentage for 30 & 200 ppm nickel. SEM images revealed that AC was a microparticle with 25 μm as an average size, while silica were nanoparticles having an average size of 12 nm. Silica/AC (2:3) composite was the most effective microparticle for nickel disposal and it is highly recommended to be used in water treatment for its high adsorptive capacity followed by AC and silica nanoparticles.
Aim of the work
The presence of toxic heavy metals in water has caused several health problems with animals, plants, and human. So that the removal of toxic heavy metals from polluted waters are one of the most important issues of environmental remediation.
The development of new products which are abundant in nature, cheap and have no environmental impact for treatment of natural resources is an important area of material technology. Calcium alginate and its composites fulfills both characteristics and have the ability to eliminate heavy metals from industrial streams.
Hence the aim of the present work is to synthesize calcium alginate and different calcium alginate composites from clays and biowaste materials (egg shell and bovine bones) and their characterization using XRD, FTIR, EDX and SEM.
The second aim of this research is to use the prepared powder samples to remove two of most hazardous heavy metals (Pb2+ and Fe3+) and measure the efficiency of each sample for remediation process.
2. Experimental and Methods
2.1. Materials, Solutions and Chemicals:
Two different biowaste materials are used in preparation of Nano hydroxyapatite –calcium alginate composites are listed in Table (3).
Table (3): Raw biowaste materials and their sources.
Material Source
Egg shell Local Hen’s egg shells
Bovine bone Local Butcher shop (bovine femur bone)
High purity analytical grade chemical material have been used in the current study are listed in Table (4).
Table (4): Solutions and chemical materials used in the current study .
Materials Chemical composition Manufacture
Diammonium hydrogen phosphate (NH4)2HPO4 Oxford laboratory India
Ammonia solution NH4OH BDH Analar England
Lead nitrate Pb(NO3)2 BDH
Iron nitrate Fe(NO3)2.9H2O BDH
Calcium chloride CaCl2 ANALAR
NaOH solution NaOH WINLAP
HNO3 solution HNO3 WINLAP
Sodium alginate C6H7O6Na Oxford laboratory India
Gelatin - Oxford laboratory India
Raw material powders are prepared and their designation in the present study is given in Table (5).
Table (5): Designation of Raw materials source in the present study.
Raw material The Source
HAP (1) Egg shells calcined at 900 °C
HAP (2) Bovine bone calcined at 1000 °C
Bentonite Abu Zaabal Fertilizer & Chemicals Co.( originating from china )
Metakaoline Kaolin from Sinai Peninsula calcined at 800 °C
Diatomeous Kazakhstan
2.2 Preparation of calcium alginate composite powders:
2.2.1 Calcium alginate:
Calcium alginate powder is prepared by using controlled gellification method [171] with some modification reported by Daemi and Barikani [172]. CA Nanoparticles are obtained by addition of CaCl2 (0.05 M) to solution containing sodium alginate (3%) by mechanical stirrer at high stirring rates (Fig. 10). Six gram of polysaccharide is dissolved in 200 mL of deionized water with high rate stirring at room temperature for 1 h. After homogenization of sodium alginate solution by mechanical stirrer, the solution (1 litre) of 0.05 molar calcium chloride is added to the system. After 1 h of the rotation, it permitted to stand at room temperature for 24 hrs. prepared nanoparticles are purified by centrifugation for 30 min. The precipitate is washed and filtered three times using double distilled water to remove the adsorbed sodium and chloride ions. The filtered CA precipitate is dried at 60 °C for 12 hours in a dry oven. This dried solids is finely grinded and sieved below 63 μm before characterization and usage.
Figure (10): preparation of calcium alginate nanoparticles
2.2.2. Bentonite:
Bentonite of Abu Zaabal Fertilizer & Chemicals Company which originating from china (Table 5) is used as a raw material in the preparation of CA – Bentonite composite. Natural bentonite dried in the oven with a temperature of 80oC for 12 hrs with the aim to eliminate moisture, bentonite ground with mortal to break chunks of bentonite then calcinied at a temperature of 800oC for 2 hours which aims to eliminate Cl bond on bentonite, and the results in the form of nanoparticles of bentonite [173].
2.2.3. MetaKaolin:
Kaolin from Sinai Peninsula is used as a raw material in the current study. Kaolin is calcined at 800 oC for 2 hrs to obtain Metakaolin (Table 5) that used in preparation of CA – Metakaolin composite.
2.2.4. Diatomeous silica:
Diatomeous which originating from Kazakhastan (Table 5) is used as a raw material in the preparation of CA – Diatomeous composite. Diatomeous is dried at temperature of 80 oC for 12 hrs with the aim to eliminate moisture.
2.2.5. Hydroxy apatite:
HAP (I)
The Egg shells mainly contain calcium carbonate (91% - 94%), calcium phosphate (1%) and other organic matters, which makes it preferable for synthesizing CaO [174]. The Nano-hydroxyapatite HAP (I) is prepared from hen’s egg shells by a method described by Laonapakul [175] with little modifications. About 50 gm. of egg shells is boiled for 30 minutes in hot water and the protein membrane is removed manually. Egg shells are dried at 80 ºC for 6 hrs. Dried eggshells solid are grinded in the agate mortar into a fine powder. The fine eggshells powder is calcined at 900 ºC for two hours in order to remove any organic residue. At this temperature the eggshells convert into calcium oxide (CaO), according to the following reaction:
Ca CO3 + Heat → CaO + CO2↑
Calculating the stoichiometric of Ca / P molar ratio = 1: 0.67 solution is prepared from CaO and DAP. Firstly, about 8.4 gram of CaO is dissolved in 100 ml 2M HNO3 and then 11.885 gm of DAP is added dropwise to calcium solution while stirring and maintaining a stoichiometry of Ca/P ratio of 1.67. NH4OH solution is added dropwise to the mixture. The pH of the solution is maintained at pH=10. A white precipitate solution is obtained and vigorously stirred for 30 minutes and permitted to stand at room temperature for 24 hours. The left white precipitate is filtered and washed three times using double distilled water to remove the adsorbed ammonia and nitrate ions. In order to obtain the final HAP solid, the filtered hydroxyapatite precipitate cake is dried at 80 °C for 10 hours in a dry oven. This dried powder is heated at 700 °C for 2 hrs. in air using control electric muffle furnace, employing a heating rate of 10 °C/min. Afterwards, the calcined powder is finely grinded and sieved below 63 μm before characterization and usage.
HAP (II)
One of the raw materials in the present study is bovine femur bones obtained from local butcher market. HAP obtained from bovine bones is prepared by a method conducted by Agnieszka et al. [176] with little modifications .In the beginning, the bovine bones (about 150 gm) are crushed into small pieces (1-2 cm) and then boiled in 1 M NaOH for one hour then in hot water for 1.5 hrs. for defatting and easier removal of the organic residues and macroscopic adhering impurities. The bones are washed and cleaned well with water and for several times afterwards. The process is followed by drying the bones at 80°C for 6 hrs. to evaporate the adsorbed water. The solid bone pieces is calcined at 1000°C for 2 hrs. at heating rate 10°C /min. and afterwards are cooled slowly to room temperature. The final solid is grinded and sieved below 63μm, and kept for characterization and usage.
2.2.6. Calcium alginate - Composites (CACS):
The compsites are prepared by mixing the CA with approprate amounts of additaves (Bentonite, Metakaolin, Gelatin, HAP(1), HAP(2) and Diatomeous) with 2:1 ratio, respectively. 6 gram of polysaccharide (sodium alginate) is dissolved in 200 mL of deionized water with high rate stirring at room temperature for 1 hrs. After homogenization of sodium alginate solution by mechanical stirrer, three grams of the chosen additives is add, then 1 litre of 0.05M calcium chloride solution is added. After 1 hour of the rotation, it permitted to stand at room temperature for 24 hrs. The prepared nanoparticles are purified by centrifugation for 30 min. The precipitate is washed and filtered three times using double distilled water to remove the adsorbed sodium and chloride ions. The CA- composites precipitates are dried at 80 °C for 24 hours in a dry oven. This dried solid are finely grinded and sieved below 63 μm before characterization and usage. Table (6) shows The synthesized composites and its abbreviations.
Table (6): The synthesized composite:
Composite Compound Abbreviations
1 Calcium alginate-Bentonite CAB
2 Calcium alginate-Metakaolin CAMK
3 Calcium alginate-HAP(1) CAHA(I)
4 Calcium alginate-HAP(2) CAHA(II)
5 Calcium alginate-Diatomeous CAD
6 Calcium alginate-Gelatin CAG
2.3. Chemical analysis of raw materials:
A- Diatomeous Silica:
Table (7): Chemical composition of Diatomeous used in the present study (XRF fused bed)
Compound Wt. %
SiO2 71.50
Al2O3 10.40
CaO 0.74
Fe2O3 3.66
MgO 1.23
SO3-- 0.71
Na2O 0.78
K2O 1.25
Cl- 0.28
TiO2 1.04
P2O5 0.10
Mn2O3 0.02
Total 91.71
Loss on ignition 8.10
Total 99.80
B- Bentonite:
Table 8: Chemical composition of bentonite used in the present study (XRF fused bed)
Compound Wt. %
SiO2 55.11
Al2O3 17.27
CaO 0.99
Fe2O3 9.03
MgO 2.27
SO3-- 0.34
Na2O 3.67
K2O 1.19
Cl- 0.62
Total 90.49
Loss on ignition 9.42
Total 99.91
C- kaolin:
Table (9): Chemical composition of kaolin used in the present study (XRF fused bed)
Compound Wt. %
SiO2 47
Al2O3 37
CaO 0.20
Fe2O3 0.20
MgO 0.02
Na2O 0.15
K2O 0.04
TiO2 1.30
Total 85.91
Loss on ignition 13.40
Total 99.31
2.4. Preparation of heavy metal ion solutions:
The metal cations of lead (Pb2+) and Iron (Fe3+) are used in the present study in the form of salts: Pb(NO3)2 and Fe(NO3)3.9H2O. One liter stock solution of each metal cation is prepared using double distilled water with metal ion concentrations of 100 ppm.
2.5. Heavy metals uptake reaction:
The metal cations reactions are conducted as follow: 20 mg of each calcium alginate composite solids is equilibrated for different time periods (5, 10, 15, 20, 30, 60 and 120 min. respectively) in glass vials with 10 ml metal cation solution with continuous shaking. After different time intervals, the solid phases are separated by centrifugation, the supernatant solution was collected for chemical analysis using Perkin Elmer 2380 Atomic Absorption Spectro Photometer (Figure 11).
Fig. 11: Perkin Elmer 2380 Atomic Absorption SpectroPhotometer.
2.6. characterization of Composite solids:
The prepared Composite solids are characterized by using Scanning Electron Microscope (SEM), X-ray diffraction (XRD), Fourier Transform Infrared (FTIR) spectroscopy and Energy Dispersive Analysis X-ray (EDX) techniques.
2.6.1. X-ray diffraction (XRD):
An X-ray diffractometer (Philips X’ PERTMPD, America, with Cu Kα radiation, 40KV and 30mA) is used to determine the mineral phases and crystallinity of the different composite powders; (Figure12).
Fig. 12: Philips X-Ray Diffractometer
The specification criteria of XRD are adjusted at 2Ө range = 5° - 60° and λ = 1.54 Ǻ at a scanning speed of 2° /min. Each composite is used to fill the aluminum mold of the diffractometer with an average thickness of about 10 mm. The obtained phases are identified by correlation with the corresponding joint committee on powder diffraction standard card (JCPDS). The average crystallite size (D) of the obtained composite powders is calculated from XRD using the Scherrer formula [177] as shown below.
where:
λ= the wave length of the X-ray.
β = the full width at half maximum (FWHM) of the peak at the maximum intensity
Ө = the diffraction angle
2.6.2. Attenuated Total Reflection Fourier Transform Infrared (ATR-FTIR):
(ATR-FTIR) technique is used to determine the main constituent chemical functional groups of the different prepared samples and the type of chemical bonding between the different atoms existing in the groups. ATR-FTIR spectrometer (Bruker, Germany Alpha-p) is configured with ATR-FTIR sample cell including a diamond crystal with a scanning depth up to 2μm. Sample powders are applied to the surface of the crystal then locked in placed with a”clutch – type” lever before measuring. The excitation of the corresponding elections when subjected to IR radiation is reflected in the spectrum as absorption bands at wavelength range from 4000 - 400 cm-1 at scanning speed of 2 cm-1 (Figure 13).
Fig. 13: ATR-Fourier Transform Infrared Spectrophotometer
2.6.3. Scanning Electron Microscope (SEM) and Energy Dispersive Analysis
X-ray (EDX):
Scanning Electron Microscopy (SEM) Model Quanta 250 FEG (Field Emission Gun) attached with EDX Unit (Energy Dispersive X-ray Analyses), with accelerating voltage 30 KV, a magnification of 14x up to 1000000, Gun.1n. FEI Company, Netherlands (Figure 14) is used to examine the surface morphology of the different prepared composite powders. The investigated samples are coated with gold (conductive layer) before imaging using EMITECH K550k sputter coater England. EDX analyzer is used to detect the chemical composition of the synthesized powders. The EDX system has a super ultra-thin window which means that it can analyze a wide range of elements.
Fig. 14: Scanning Electron Microscope with EDX
2.7. Adsorption studies
An accurately 0.02 g of CACs is added into 10 mL of solutions in a 25.0 mL glass tubes containing the specified concentrations of metal ions. The mixtures are shaken at 25.0 oC for a fixed period (2 hrs.) and at the end of shaking periods, the contents are filtered through filter paper. The filtrate is analyzed for final metal concentration using Perkin Elmer 2380 Atomic Absorption Spectro Photometer. Each experiment is performed in triplicate and the average of the results is recorded.
The effects of contact time (5-120 min.) and metal ion concentration on the sorption process are realized using the same methodology. The amounts of metal ion adsorbed onto CACs sorbent, qe (mg g–1), are calculated using the Equation:
qe = (Co – Ce) V/ m
where Co and Ce (mg L–1) are initial and equilibrium concentrations of metal ions, m (g) is the weight of sorbent in the solution and V (L) is the volume of the solution.
The efficiency of adsorption (removal %) is calculated according to the Equation:
% Removal = [(Co – Ce) /Co] × 100
2.7.1. Effect of contact time:
To study the effect of contact time on the adsorption efficiency of Pb2+ and Fe3+ ions by CACs, 0.02 g of CACs is added into 10 mL of 100 mg L-1 Pb2+ (pH 5.7 and 4) and Fe3+ ( pH 2.6) solutions at 25±1°C with time interval from 5 -120 minutes.
2.7.2. Effect of pH:
The effect of pH on the adsorption is performed only for Pb2+. The study is achieved with two pHs values (4 and 5.7) the original solution of lead is at pH 5.7 and to get pH 4 we use HNO3 solution (0.01 M). This process is conducted at 30 minutes contact time, 20 mg dosage of the different composites and 100 ppm of metal solution at 25±1°C. With respect to Fe3+ solution, the experimental work is carried out at pH =2.7 only. This is due to the precipitation of Fe3+ at pH ≥ 3.
2.7.3. Effect of dosage:
To study the effect of dosage (10, 15, 20 mg) on the adsorption efficiency of Pb2+ on CA-Np (as a standard model). The previous dosages are added to 10 mL of 100 mg L-1 Pb2+ at 25±1°C with 30 minutes contact time.
2.8. Adsorption kinetics:
Adsorption kinetic studies are important since they describe the solute uptake rate which controls the residence time of adsorbate at the solid–liquid interface and also provide valuable insights into the reaction pathways.
Pseudo–first–order (Equation I) and pseudo–second–order (Equation II) models were applied in order to investigate the adsorption kinetics of Pb2+ and Fe3+ ions onto CACs. The conformity between experimental data and the model-predicted values is expressed by the correlation coefficients (R2). Meanwhile, the capacity values calculated from the pseudo–first and second–order models are compared with that obtained from the experimental data. The kinetic models can be presented as follows,
ln (qe –qt ) = ln qe – k1 t (I)
t/qt = 1/k2qe2 + t/qe (II)
where qt is the amount of metal ion adsorbed (mg g-1) at time (t), qe is the maximum adsorption capacity (at equilibrium) (mg g-1) and k1(min−1) is the rate constant for pseudo–first–order sorption, qe is the maximum adsorption capacity (at equilibrium) (mg g-1) and k2 (g mg-1 min-1) is the rate constant for the second-order sorption. from plots of t/qt versus t for the second–order reactions, the k2 and qe values are calculated by using the values of intercept and slope.
2.9. Adsorption isotherm models:
The adsorption equilibrium is investigated by using two well-known isotherm models (Langmuir and Freundlich) to provide the fundamental physicochemical data and investigate the applicability of sorption process at a fixed temperature. The equilibrium conditions of the adsorption process are described by utilizing the linearized equations indicated below,
Ce/qe =Ce/qm + 1/ KL qm
where qm (mg g–1) and KL (L mg–1) are constants in Langmuir’s equation which are referred to the maximum adsorption capacity and the Langmuir model constant that is indirectly related to the energy of adsorption. Also qe and Ce parameters represent the equilibrium adsorption capacity and the equilibrium concentration heavy metal ion, respectively.
Freundlich adsorption isotherms assumes aheterogeneous surface with anon-uniform distribution of heat of adsorption over the surface with the possibility of the multilayer adsorption [178, 179].the Freundlich equation which is expressed as
qe =log KFCe1/n
where KF is the measure of adsorption capacity and n is the adsorption intensity linear form of Freundlich
log qe =log KF + (1/n) log Ce
where qe is the amount adsorbed (mg/g), Ce is the equilibrium concentration of adsorbate (mg/L), and KF and n are the Freundlich constants related to the adsorption capacity and adsorption intensity, respectively. A plot of log qe vs. log gives a linear trace with a slope of 1/n and intercept of log KF.
3. Results and Discussion
3.1. characterization of the synthesized composites.
3.1.1. Fourier Transform Infrared (FTIR) analysis:
The Fourier transform infrared (FTIR) spectra of calcium alginate and calcium alginate composites are given in Figures (15-18). Also their frequencies of the absorbtion bands of FTIR spectra and their structure assignments are given in Table (10). Spectrum of calcium alginate (Figure 15), showed important absorption bands regarding hydroxyl, ether and carboxylic functional groups. Stretching vibrations of O–H bonds of alginate appeared in the range of 3000–3600 cm-1 particualy at 3444 cm-1. Stretching vibrations of aliphatic C–H are observed at 2920–2850 cm-1. The observed bands at 1632 and 1454 cm-1 are attributed to asymmetric and symmetric stretching vibrations of carboxylate salt ion, respectively. 1154 cm-1 (CO-stretching of ether group) and 1025 cm-1 (C-O stretching of alcohol group)[180].
The spectra of CAG (Fig.16), the absorption band at around 3442 cm-1 concerned with OH stretching vibration for CA slightly broadened and shifted to a lower wave number with the blending with gelatin, suggesting the formation of an intermolecular hydrogen bond [181]. The strong absorption band at 1631 cm-1 for CA assigned to the asymmetric stretching vibration of COO- has coupled with the absorption band at 1631 cm-1 in gelatin.
CA raw material, reveals asimilar spectra with CAB as shown in Fig 17, in addition some peaks are found at 3699 and 3444 cm-1 are due to lattice OH and bound water stretching vibrations. A strong and sharp band is detected at 1022 cm-1 which is related to Si–OH stretching vibrations [182]. Peaks found at 1384 cm-1 is due to CO3 stretching of calcite, 1034 cm-1 assigned to Si-O stretching, and 875 cm-1 is due to OH bending of the Al-Al-OH group. A similar OH bending vibration is observed for Al-Mg-OH at 842 cm-1, 690 cm-1 assigned to quartz. Also, there is a shoulder peak at 520 cm-1 (Al-O-Si bending), and 464 cm-1 (Si-O-Si bending)[183].
The IR spectra of CAMK and CAD composites revealed a similar spectra with CAB composite spectra (Figure 17). CAD showed a strong band at 1084 and 1048 cm-1 due to Si–OH and Si-O vibrations. These bands overlabed with the band at 1025 cm-1 due to CO-stretching of alcohol group in alginate. A strong and sharp peak at 470 cm-1 are also detected due to Si-O-Si bending. The band of Si-O-Si bending in CAD is the strongest and sharpest compared to CAMK and CAB.
The HAP-Alginate samples (CAHA(I) and CAHA(II)) (Figure 18) revealed a similar spectra, at 1625-1630 cm-1 and 3440-3445 cm-1 (due to the presence of free water), 1454 and 874 cm-1 ( due to CO32- ions), 3570 and 630-633 cm-1 (due to structural OH of hydroxy apatite). These peaks due to the hydroxyapatite phase [184, 185]. The most intensive bands in the range of 1044 –1090 cm-1 corresponded to the triply degenerated asymmetric stretching vibrations of P-O. Otherwise the peak at 962.97 cm-1 indicates the non-degenerated asymmetric mode of PO43-. The very strong and sharp bands observed at 569-572 and 602 -603 cm-1 attributed to triply degenerated bending mode of the O-P-O in PO43- group. The larges parting distance of these bands revels the crystalline phase [184, 185].
Table (10): Assignments of the absorption bands of the IR spectra λ (cm-1) of the prepared composites
Peak Assignment Strength CA CAB CAMK CAHA(I) CAHA(II) CAD CAG
Structural OH of addittives w - 3699 - 3571 3570 - -
OH stretching mode of adsorbed water molec. or OH of alginate lattice structure s , b 3444 3444 3444 3444 3444 3443 3442
Stretching mode of aliphatic C–H. w 2924 2923 2924 2923 2924 2924 2924
asymmetric stretching vibrations of carboxylate salt ion m 1631 1634 1631 1629 1630 1629 1631
symmetric stretching vibrations of carboxylate salt ion vw 1427 1461 1431 1433 1419 1433 1427
CO-stretching of ether group vw 1155 1150 1153 - - 1103 1155
CO-stretching of alcohol group w 1024 1033 - - - - 1024
Si–OH stretching vibrations mode w - 1080 1081 - - 1084
CO3 stretching of calcite w 1384 1384 1384 1384 1384 1384 1384
Si-O stretching w - 1033 1052 - - 1048 -
OH bending of the Al-Al-OH group vw - 875 850 - - 845 -
OH bending vibration of Al-Mg-OH group of quartz vw - 842
690 777 - - 797 -
Al-O-Si bending vw - 520 510 - - 526 -
bending vibration mode of Si-O-Si s - 464 456 - - 470 -
vibration mode CO32- vw - - - 1454
874 1457
874 - -
structural OH in HA w - - - 631 631 - -
triply degenerated asymmetric stretching vibrations of P-O vs - - - 1044
1090 1048
1089 - -
non-degenerated asymmetric of PO43- vw - - - 962 962 - -
triply degenerated bending mode of PO43- ms - - - 602
569 602
571 - -
s = strong w = weak vw = very weak b = broad m = meadium
Figure 15: FTIR Spectra of Calcium alginate.
Figure 16: FTIR Spectra a- CA, b- CAG
Figure 17: FTIR Spectra a- CA, b- CAB, c- CAMK ,d- CAD.
Figure 18: FTIR Spectra a- CAHA(I) , b-CAHA(II)
3.1.2. X-ray diffraction for crystal phase detection:
Figure(19) presents the X-ray diffraction pattern of CA and CA composites in the range of 2θ = 5-60o diffraction degree . Two typical peaks in 2θ =16° and 22° are observed for calcium alginate. The XRD of CAG shows typical peaks around 12° and 21°[186].
The samples of CAB, CAMK and CAD powders revealed very similar XRD pattern peaks of quartz phase. These composites showed the maximum relative intensity (I/I0) peak of 100% quartz in the rang of 2θ =26.65o and 20.85o and d (Ao) value spacing equal 3.34 and 2.8, (reference code :01-070-3755). On the other hand, additional peaks are detected due to the presence of minor amounts of calcite (CaCO3) and halite (NaCl). The presence of calcite phase may be attributed to carbonation of calcium during composite synthesis, while the presence of NaCl may be due to the reaction of Na+ ions of alginate with Cl- ions of CaCl2 solution and its trapping bettween the layers, which can not completely removed by washing.
The well resolved XRD peaks of The sample HAP(I) ,CAHA(I), HAP(II) and CAHA(II) could be easily indexed on the basis of hexagonal crystal system of space group P63/m with respect to JCPDS file no. 9-432. They also revealed very similar XRD pattern peaks of 100 % HAP (reference code :01-086-1194) in the rang of 2θ =31o and 32o and d (Ao) value spacing equal 2.81 and 2.78. There is no any considerable shifts in 2θ are detected between HAP and HAP composite . The diffraction peaks of HA and CA-HAP exhibit sharp diffraction peaks which indicate the high crystallinity of the structure and there is no any additional phases are detected.
XRD analysis of HAP(II) and CAHA(II) also indicated the absence of secondary phases, such as tri calcium phosphate (TCP) or calcium oxide (CaO). In the case of HAP(I) and/or CAHA(I), their diffraction patterns revealed additional phase of β-tricalcium phosphate, beside the HAP as a main phase.
Hydoxyapatite prepared from bovine bone contain certain amount of carbonate (CO32-) in its lattice structure [175]. So the sample CAHA(I) are expected to be a carbonated apatite type, and carbonate ions affect on the the degree of crystallinite. For this reason, CAHA(II) ( HA prepared from bovine bones) revealed higher crystal size (83 nm) than CAHA(I) which prepared from eggshells (59 nm). There is no significant differences in the crystal sizes of the prepared HAP and HAP Compsites. Table (11) showed the crystal size, 2θ, d-spacing(oA), maximum relative intensity (I/Io) peak and the main phase detected for the prepared adsorbents
Table (11): 2θ, d-spacing(oA), Crystal size ( nm ) of the maximum relative intensity (I/Io) peak and the main phase detected for the prepared adsorbents.
Character
2θ
d-spacing (Ao)
Crystal size (nm)
Main phase detected
CAHA(I)
31.83
2.81
59
HA
CAHA(II)
31.77
2.81
83
HA
CAB
26.65
3.34
83
Quartz
CAMK
26.68
3.34
84
Quartz
CAD
26.63
3.34
59
Quartz
Figure 19: X-ray diffraction pattern of the prepared calcium alginate composite powders.
Figure 19: continue
Figure 19: continue
Figure 19: continue
3.1.3. Scanning Electron Microscope (SEM) and Energy Dispersive X-ray (EDX)
The SEM images and EDX elemental analysis of synthesized calcium alginate and calcium alginate composites are shown in Figure (20). The images are taken at 300× and 2500× magnification. SEM investigations showed the presence of agglomerates of irregular shape particles.
According to the results obtained by EDX analysis, carbon, oxygen and calcium are the major constituents of all the prepared composites related to calcium alginate polymer present, Also, Almost prepared composites confirmed the presence of calcite and halite as additional phases in minor amounts which is in aggreement with X-ray results.
EDX analyses of CA and CAG revealed that the major elements present are carbon and oxygen corresponding to polymer composition. The SEM of CAG (Fig.20(i)) showed a smooth and homogeneous morphology, suggesting high miscibility and blend homogeneity between calcium alginate and gelatin.
In the case of CAB, CAMK and CAD, the presence of an obvious peak related to the Si compounds is evident. According to the results obtained by EDX analysis, weigh percent of elements present indicating a large portion of the composites is composed of Si compounds which is suitable for an efficient sequestering metal cations from aqueous solution[129].
EDX analyses of the prepared HAP revealed that inorganic phases of bovin bone and egg shells were mainly composed of calcium and phosphorus as the major constituents with some minor components such as C, O, Na, Mg and Si. The weight and atomic percentage shows that the Ca/P ratio around 1.7 and 1.8 which is below 2 and acceptable where the ideal Ca/P ratio of HA is 1.67[174], in HAP composites (CAHA(I) and CAHA(II)) this ratio increased than 2 due to the excess amount of Ca crosslinkage of calcim alginate presents.
a) CA
Element Wt % At %
C K 35.08 46.27
O K 45.62 45.17
Na K 2.08 1.43
Cl K 6.18 2.76
Ca K 11.05 4.37
Total 100 100
b) CAB
Element Wt % At %
C 16.29 25.90
O 37.59 44.87
Na 1.14 0.95
Al 13.61 9.77
Si 16.26 11.06
Cl 6.83 3.68
Ca 7.08 3.37
Ti 1 0.4
total 100 100
Figure 20: SEM images and EDX analysis of prepared samples a) CA & b) CAB .
c) CAMK
Element Wt % At %
C K 9.75 16.10
O K 41.60 51.55
Na K 1.07 0.93
Al K 4.43 3.26
Si K 26.61 20.91
Cl K 6.25 3.50
Ca K 5.85 2.90
K K 0.98 0.49
Mg 0.45 0.37
Total 100 100
d) CAD
Element Wt % At %
CK 27 39.27
OK 39.73 43.39
Na K 2.46 1.87
Al K 2.30 1.49
Si K 5.16 3.21
Cl K 7.79 3.84
Ca K 14.62 0.21
K K 0.46 0.21
Mg 0.47 0.34
Total 100 100
Figure 20: SEM images and EDX analysis of prepared samples c) CAMK & d) CAD.
e) HAP(I)
Element Wt % At %
C K 4.66 9.69
O K 29.05 45.31
P K 20.45 16.47
Ca K 45.84 28.53
Total 100 100
f) HAP(II)
Element Wt % At %
C K 3.96 8.22
O K 28.88 45.05
Na K 2.42 2.62
Mg K 0.76 0.78
P K 18.95 15.27
Ca K 45.03 28.04
Total 100 100
Figure 20: SEM images and EDX analysis of prepared samples e) HAP(1) & f) HAP(2).
g) CAHA(I)
Element Wt % At %
C K 20.69 33.23
O K 36.42 43.90
Na K 1.74 1.46
P K 10.18 6.34
Cl K 2.62 1.42
Ca K 28.35 13.64
Total 100 100
h) CAHA(II)
Element Wt % At %
C K 7.32 14.25
O K 31.95 46.71
Na K 0.70 0.71
Mg k 0.40 0.39
P K 17.06 12.86
Cl K 2.71 1.79
Ca K 39.86 23.27
Total 100 100
Figure 20: SEM images and EDX analysis of prepared samples. g) CAHA(I) & h) CAHA(II).
i) CAG
Element Wt % At %
C K 39.21 49.41
O K 44342 44.07
Na K 2.08 1.43
Cl K 3.15 1.8
Ca K 11.05 4.37
Total 100 100
Figure 20: SEM images and EDX analysis of prepared samples. i) CAG
3.2. characterization of the composites after metal ion uptake:
3.2.1. FTIR Analysis after metal ion uptake:
FTIR spectra of CA and CAHA(I) after targeted metal ions uptake ( Pb2+, Fe3+ ) are shown in Figures 21 and 22 respectively and their peak assignments are represented in Table (12). The result of FTIR spectra showed that there is no new absorption peaks were detected . There are alittle peak shifts which may be attributed to the corporation and substitution of metal ion in the lattice structure of CA and CAHA(I).
Table 12: Change in the absorption bands of the IR spectra λ (cm-1) of CA and CAHA(I) powder after metal ion uptake
Peak Assignment Strength CA
CA
+
Pb2+
CA
+
Fe3+ CAHA(I)
CAHA(I)
+
Pb2+
CAHA(I)
+
Fe3+
Lattice Structural OH of hydroxyapatite w -
- - 3571 3571 3571
OH stretching mode of adsorbed water molec. or OH of alginate lattice structure s , b 3444 3444 3434 3444 3444 3444
Stretching mode of aliphatic C–H . w 2924 2924 2923 2923 2924 2923
asymmetric stretching vibrations of m 1631 1598 1631 1629 1599 1615
symmetric stretching of COO- ion vw 1427 1425 1424 1433 1433 1433
CO-stretching of ether group. vw 1155 1156 1155 -
CO-stretching of alcohol group w 1024 1021 1021 -
CO3 stretching of calcite w 1384 1384 1383 1384 1384 1381
vibration mode CO32- vw - 1454
874 1454
874 1458
874
structural OH in HA w - 631 631 631
triply degenerated asymmetric stretching vibrations of P-O vs - 1044
1090 1045
1090 1045
1090
non-degenerated asymmetric of PO43- vw - 962 962 962
triply degenerated bending mode of PO43- ms - 602
569 602
570 601
568
s = strong w = weak vw = very weak b = broad m = meadium
Figure 21: FTIR Spectra a- CA , b-CA+ Pb2+ , c- CA + Fe3+
Figure 22: FTIR Spectra of a- CAHA(I) , b- CAHA(I)+ Pb2+ , c- CAHA(I)+ Fe3+
3.2.2. X-ray diffraction for crystal phase detection after metal ion uptake:
XRD patterns of CA and CAHA(I) after Pb2+ and Fe3+ metal ions removal did not revealed any new phases. supported the proposal that Pb2+ and Fe3+ ions uptake was not dependent on dissolution/precipitation mechanisms. Pb2+ and Fe3+ ions removal may be occurs by adsorption mechanisms like surface complexation or ionic exchange [187]. As it can be seen in Figures 23-24, XRD patterns showed some changes in their relative intensities and crystal sizes (Table 13). Also, ther are some little shifts in d- spacing values, this may be due to the ion exchange between Ca2+ and metal cations of Pb2+ and Fe3+ in lattice structure of CA and CAHA(I) nanopowders.
Table 13: 2θ, d-spacing(Ao), Crystal size ( nm ) of the maximum relative intensity (I/Io) peak of CAHA(I) before and after metal ion uptake.
Character
2θ d-spacing (Ao) Crystal size (nm)
CAHA(I)
31.83 2.811 59
CAHA(I) + Pb2+
31.74 2.819
60
CAHA(I) + Fe3+ 31.71 2.821 37
Figure 23: X-ray diffraction pattern of CA after metal ion uptake
Figure 24: X-ray diffraction pattern of CAHA(I) after metal ion uptake
3.2.3. Scanning Electron Microscope (SEM) and Energy Dispersive X-ray (EDX) after metal ion uptake:
SEM of CA and CAHA(I) after metal ion uptake revealed some changes in morphology and microstructure(Fig. 26,27). Morever EDX results indicated the presence of Pb2+ and Fe3+ ions with CA and CAHA(I). As can seen in Figures 26 and 27, the peak of Fe3+ is more intense than that of Pb2+ and this confirms, the higher value of Fe3+ ions uptake by CA and CAHA(I) comparing with Pb2+ ions. The weight and atomic percentage of Ca2+ ions in CA and CAHA(I) after metal ion uptake was less than its value before metal ion removal as listed in Table (14). The decrease in calcium percentage may be attributed to ion exchange between targeted metal ions and; (1)calcium ions of calcium alginate in CA and CAHA(I) as follows:
Ca(ALG)2 + Pb2+ Pb(ALG)2 + Ca2+
(2) or Ca ions of HA present in CAHA(I). This ion exchange mechanism between Pb2+ ions (as example) and Ca2+ ions of HA produced anew phase of hydroxypyromorphite [132, 188], this mechanism is expressed as:
Ca10 (PO4)6 (OH)2 + x Pb2+ x Ca2+ + Ca10-xPbx (PO4)6(OH)2
However , this phase is not detected in the present study this may be attributed to under limit of XRD
Table 14: Ca2+ ion percentage in CA and CAHA(I) before and after metal ion uptake.
Character
Ca2+ % before uptake
Ca2+ % after uptake
Wt %
+Pb2+ (Wt %)
+ Fe3+ (Wt %)
CA 11.05 5.69 1
CAHA(I) 28.35 15.61 23.43
Also the uptake process may be occurs during chelation bonding of targeted metal ions with two carboxylic groups of alginate and one or two OH sites of the alginate ring (Fig 25) [189]. In this case metal ions may forms complexes with two adjacent alginate rings. Here,‘‘adjacent’’ means either two neighbor alginate rings of a single polymeric chain (intramolecular chelation) or two rings from two parallel chains (intermolecular chelation) [189].
Figure 25: potential active sites of CA which may bonds with targeted metal ion
a) CA + Fe+3
Element Wt % At %
C K 15.77 32.25
O K 27.57 42.32
Si K 0.41 0.35
Cl K 0.62 0.43
Ca K 1 0.61
Fe K 54.64 24.03
Total 100 100
b) CAHA(I) + Fe+3
Element Wt
% At
%
C K 7.60 12.77
O K 36.54 55.02
Al k 0.37 0.33
P K 14.03 12.91
Ca K 23.43 11.52
Fe K 18.03 7.78
Total 100 100
Figure 26: SEM images and EDX analysis of CA and CAHA(I) after iron removal
a) CA + Pb+2
Element Wt % At
%
C K 32.44 51.66
O K 36.01 43.05
Al k 0.30 0.21
Pb M 25.56 2.36
Ca K 5.69 2.71
Total 100 100
b) CAHA(I) + Pb+2
Element Wt % At %
C K 9.53 18.79
O K 32.94 48.74
Al k 0.37 0.32
P K 15.46 11.81
Pb M 8.99 1.03
Ca K 32.71 19.31
Total 100 100
Figure 27: SEM images and EDX analysis of CA and CAHA(I) after lead removal.
3.3. Metal ion uptake by adsorption process.
3.3.1. Effect of contact time on the adsorption process.
3.3.1.1 Calcium Alginate (CA):
The effect of contact time on the adsorption capacity of CA for Pb2+ ( natural pH of 5.7), Pb2+ (pH 4) and Fe3+ ( natural pH of 2.6) is represented in Figure 28. The equilibrium points of adsorption were attained within the first 60 min. for Pb2+ (82.3%) , 30 min. for Pb2+ (86.82%) and 30 min. for Fe3+ (94.15%) of contact time. the adsorption of Pb2+, Pb2+( pH 4) and Fe3+onto CA was rapid in the initial stages up to 30 min, and was almost same at high contact time. This behavior could be attributed to the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption process. The exponential phase for Pb2+( pH 4) and Fe3+ was found to be between 5 minutes and 30 minutes. For Pb2+ it was found to be between 5 minutes and 60 minutes. Further increase in contact time led to no significant adsorption of metal ions by CA probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches its maximum after 120 minutes (94.18%).
Figure 28: Effect of contact time on the adsorption of Pb2+( pH 5.7) , Pb+2(pH 4) and Fe3+ ions(pH 2.6) from aqueous solution by CA under experimental conditions of CA mass 0.02 g/10mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.1.2 CAB
The effect of contact time on the adsorption capacity of CAB for Pb2+ ( natural pH of 5.7), Pb2+(pH 4) and Fe3+ ( natural pH of 2.6) is represented in Figure 29. The equilibrium points of adsorption were attained within the first 60 minutes for Pb2+ (66.40%), 30 minutes for Pb2+ (pH 4)(68.2%) and 30 minutes for Fe3+ (89.42%) of contact time. the adsorption of Pb2+, Pb2+( pH 4) and Fe3+onto CAB was rapid in the initial stages up to 30 min, and was almost same at high contact time. This behavior may be due to the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption process. The exponential phase for Pb2+( pH 4) and Fe3+ was found to be between 5 minutes and 30 minutes. For Pb2+ it was found to be between 5 minutes and 60 minutes. The sorption efficiency by CAB was almost constant probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches its maximum after 120 minutes (89.6%).
Figure 29: Effect of contact time on the adsorption of Pb2+(pH=5.7)Pb+2(pH=4) and Fe3+(pH= 2.6)ions from aqueous solution by CAB under experimental conditions of CAB mass 0.02 g/10 mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.1.3 CAMK
The effect of contact time on the removcal of Pb2+ ( natural pH of 5.7), Pb2+ ( pH 4) and Fe3+ ( natural pH of 2.6) by CAMK is represented in Figure 30. The adsorption percent of metal ions was fast at initial stages and gradually become slower until the equilibrium is attained. The optimal contact time to attain equilibrium was experimentaly found to be about 30 min. for Pb2+ (69.26%) , Pb2+( pH 4) (77.7%) and Fe3+ (89.87%). the adsorption of Pb2+, Pb2+( pH 4) and Fe3+onto CAMK was rapid in the initial stages up to 30 minutes, and was almost same at high contact time. This behavior could be because of the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption processThe exponential phase for Pb2+, Pb2+( pH 4) and Fe3+ was found to be between 5 minutes and 30 minutes. Further increase in contact time led to no significant adsorption of metal ions by CAMK probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches its maximum after 120 minutes (90%).
Figure 30: Effect of contact time on the adsorption of Pb2+ (pH= 5.7), Pb+2( pH= 4) and Fe3+ (pH=2.6) ions from aqueous solution by CAMK under experimental conditions of CAMK mass 0.02 g/10 mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.1.4. CAHA(I)
The effect of contact time on the adsorption capacity of CAHA(I) for Pb2+ (natural pH of 5.7), Pb2+ (pH 4) and Fe3+ ( natural pH of 2.6) is represented in Figure (31). As it can be seen in fig.22, the equilibrium points of adsorption were attained within the first 60 minutes for Pb2+ (80.78%), 30 minutes for Pb2+ (pH 4) (86.62%) and 30 minutes for Fe3+ (99.33%) of contact time. the adsorption of Pb2+, Pb2+( pH 4) and Fe3+onto CAHA(I) nanopowders was rapid in the initial stages up to 30 min, and was almost same at high contact time. This behavior could be attributed to the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption process. The exponential phase for Pb2+( pH 4) and Fe3+ was found to be between 5 minutes and 30 minutes. For Pb2+ it was found to be between 5 minutes and 60 minutes. The sorption efficiency by CAHA(I) was almost constant probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches its maximum after 120 minutes (99.34%)
Figure 31: Effect of contact time on the adsorption of Pb2+ (pH= 5.7), Pb+2(pH= 4) and Fe3+ ions (pH= 2.6) from aqueous solution by CAHA(I) under experimental conditions of CAHA(I) mass 0.02 g/10 mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.1.5. CAHA (II)
The adsorption efficiency of Pb2+ ( natural pH of 5.7), Pb2+ (pH 4) and Fe3+ ( natural pH of 2.6) by using CAHA (II) was tested at different contact time. Asit can be seen in figure 32, the equilibrium points of adsorption were attained within the first 30 minutes for Pb2+ (74.08%) and Pb2+ (pH 4) (84.67%) while for Fe3+ were attained within the first 10 minutes (98.65%) of contact time. the adsorption of Pb2+, Pb2+( pH 4) and Fe3+onto CAHA(II) nanopowder was rapid in the initial stages up to 30 min, and was almost same at high contact time. This behavior could be attributed to the availability of a large number of active sites for rapid surface metal ions binding during the initial stages of adsorption process. The exponential phase for Pb2+and Pb2+( pH 4) was found to be between 5 minutes and 30 minutes. For Fe3+ it was found to be between 5 minutes and 10 minutes. Further increase in contact time led to no significant adsorption of metal ions by CAHA (II) probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches 98.78% after 60 minutes and its maximum after 120 minutes (98.81%).
Figure 32: Effect of contact time on the adsorption of Pb2+ (pH= 5.7), Pb+2(pH= 4) and Fe3+ ions (pH= 2.6) from aqueous solution by CAHA (II) under experimental conditions of CAHA (II) mass 0.02 g/10 mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.1.6. CAD
The effect of contact time on the adsorption capacity of CAD for Pb2+ ( natural pH of 5.7), Pb2+(pH 4) and Fe3+ ( natural pH of 2.6) is represented in Figure 33. The equilibrium points of adsorption were attained within the first 30 minutes for Pb2+ (68.96%), Pb2+( pH 4) (63.23%) and Fe3+ (86.72%) of contact time. The adsorption of Pb2+, Pb2+(pH 4) and Fe3+onto CAD was rapid in the initial stages up to 30 min, and was almost same at high contact time. This behavior could be attributed to the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption process. The exponential phase for Pb2+, Pb2+( pH 4) and Fe3+ was found to be between 5 minutes and 30 minutes. The sorption efficiency by CAD was almost constant probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches its maximum after 120 minutes (86.95%).
Figure 33: Effect of contact time on the adsorption of Pb2+ (pH= 5.7), Pb+2(pH= 4) and Fe3+ ions (pH= 2.6) from aqueous solution by CAD under experimental conditions of CAD mass 0.02 g/10 mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.1.7. CAG
The sorption efficiency exhibited by of CAG for Pb2+ ( natural pH of 5.7), Pb2+ (pH 4) and Fe3+ ( natural pH of 2.6) is depicted in Figure 34. The equilibrium points of adsorption were attained within the first 30 minutes for Pb2+ (83.43%), Pb2+( pH 4) (87.93%) and Fe3+ (91.11%) of contact time. The adsorption of Pb2+, Pb2+( pH 4) and Fe3+onto CAG was rapid in the initial stages up to 30 min, and was almost same at high contact time. This behavior could be due to the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption processThe exponential phase for Pb2+, Pb2+( pH 4) and Fe3+ was found to be between 5 minutes and 30 minutes.. The sorption efficiency by CAG was almost constant probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches 91.55% after 60 min. and reaches its maximum after 120 minutes (91.81%).
Figure 34: Effect of contact time on the adsorption of Pb2+ (pH= 5.7), Pb+2(pH 4) and Fe3+ ions (pH= 2.6) from aqueous solution by CAG under experimental conditions of CAG mass 0.02 g/10 mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.2. Effect of pH
The effect of pH on the adsorption is performed only for Pb2+ because of Fe3+ solution was stable only at pH lower than 3. The study is achieved with two pHs values (4 and 5.7) the original solution of lead was at pH=5.7 and pH=4 at 5-120 minutes contact time, 20 mg dosage of the different composites and 100 ppm of metal solution at 25±1°C. As can be seen in Figures 28-34, the adsorption efficiency of Pb2+ at pH=4 is higher than that of pH=5.7 for all the composites.
3.3.3. Effect of adsorbent dosage on metal ion adsorption.
The experimental results of the adsorption of Pb2+ on CA ( as astandard model ) as afunction of adsorbent dosage 10, 15 and 20 mg/10 mL, initial Pb2+ concentration of 100 mgL-1, natural pH of 5.7, temperature 25oC at the optimal contact time (30 min) and interval contact time ( 5-30 min) are shown in fig. 36 and 37 respectively. As can be seen in Figures 35 and 36, the Pb2+ adsorption percent rapidly increased with the increase in the adsorbent dosage . this can be attributed to higher adsorbent dosage due to the increased surface area providing more adsorption sites available which gave rise to higher removal of lead [190].
Figure 35: Effect of CA dosage on Pb2+ at contact time 30 min. , initial concentration of 100 mgL-1, natural pH of 5.7 and temperature 25±1oC.
Figure 36: Effect of CA dosage on Pb2+ at contact time 5-30 min. , initial concentration of 100 mgL-1, natural pH of 5.7 and temperature 25±1oC.
3.4. kinetics studies of the adsorption process.
The kinetic study is useful to predict the adsorption rate which is very important in modeling and designing of the adsorption process [191]. The pseudo-first rate equation of lagergren and pseudo-second order kinetic model, as the most widely used models, are used to evaluate the mechanism of adsorption process.
3.4.1. The pseudo first-order model:
The linear form of the pseudo first-order kinetic rate equation is given as follows:
ln (qe –qt ) = ln qe – k1 t
where qt is the amount of metal ion adsorbed (mg g-1) at time (t), qe is the maximum adsorption capacity (at equilibrium) (mg g-1) and k1(min−1) is the rate constant for pseudo–first–order sorption, qe is the maximum adsorption capacity (at equilibrium) (mg g-1). The kinetic of adsorption are evaluated at an initial concentration of 100 mg/L for Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+(pH=2.6), adsorbent dosage of 0.02 g/10 mL and temperature of 25oC. Plot of ln (qe –qt ) vs t is drawn as shown in Figures 37-43.
The rate constant at equilibrium (qe) and regression coefficient (R2) obtained from the plots of pseudo-first rate equation of adsorbed Pb2+ and Fe3+ at equilibrium (qe) for all the adsorbents are given in Tables 15,16 and 17 respectively. As it can be seen in Figures (37-43) and Tables (15-17), The regression coefficient does not close to unity. Also, the values of qe obtained from pseudo-first order equation for all the adsorbent are different and not matched notably with the experimental qe value.
Figure 37: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CA at experimental conditions of 100 mg/L metal ion concentration, CA mass 0.02 g/10 mL and 25±1oC.
Figure 38: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAB at experimental conditions of 100 mg/L metal ion concentration, CAB mass 0.02 g/10 mL and 25±1oC.
Figure 39: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAMK at experimental conditions of 100 mg/L metal ion concentration, CAAMK mass 0.02 g/10 mL and 25±1oC.
Figure 40: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+( pH= 4) and Fe3+ (pH=2.6) adsorbed on CAHA(I) at experimental conditions of 100 mg/L metal ion concentration, CAHA(I) mass 0.02 g/10 mL and 25±1oC.
Figure 41: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAHA(II) at experimental conditions of 100 mg/L metal ion concentration, CAHA(II) mass 0.02 g/10 mL and 25±1oC.
Figure 42: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+(pH=2.6) adsorbed on CAD at experimental conditions of 100 mg/L metal ion concentration, CAD mass 0.02 g/10 mL and 25±1oC.
Figure 43: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAG at experimental conditions of 100 mg/L metal ion concentration, CAG mass 0.02 g/10 mL and 25±1oC.
3.4.2. The pseudo second-order model:
The linear form of the pseudo second-order kinetic rate equation is given as follows:
t/qt = 1/k2qe2 + t/qe
where qt is the amount of metal ion adsorbed (mg g-1) at time (t), qe is the maximum adsorption capacity (at equilibrium) (mg g-1) and k2 (g mg-1 min-1) is the rate constant for the second-order sorption. from plots of t/qt versus t for the second–order reactions Figures 44-50, the k2 and qe values are calculated by using the values of intercept and slope as summarized in Tables (15, 16, 17).
The rate constant and regression coefficient (R2) obtained from the plots of pseudo-second rate equation of adsorbed Pb2+ (pH=5.7), Pb2+( pH 4) and Fe3+ (pH=2.6) at equilibrium (qe) for all the adsorbents are given in Tables 15,16 and 17 respectively . from the linear plots, the qe,experimental and the qe,calculated values are very close to each other, and also, the calculated coefficients of determination, R2, are close to unity.
Table 15: Experimental and calculated parameters of pseudo-first and second order kinetic models of Pb2+ (natural pH of 5.7) adsorbed on CA and CACs powder.
Qe experimental Pseudo first order Pseudo second order
qe calculated
K1
R2 qe calculated
K2
R2
CA
41.15
3.412
0.0769
0.9183
41.15
1.91x10-6
0.9997
CAB
33.2
5.77
0.03538
0.4826
33.69
3.4x10-5
0.9923
CAMK 34.89 11.881 0.12773 0.9636 36.16 2.41x10-5 0.9987
CAHA(I) 40.39 16.29 0.11735 0.9438 42.19 2.06x10-5 0.9984
CAHA(II)
37.09 35.55 0.19646 0.8367 39.65 3.21x10-5 0.9951
CAD 34.87 27.91 0.1423 0.8184 38.13 6.22x10-5 0.9778
CAG 42.07 11.189 0.10384 0.6573 43.3 1.33x10-5 0.9981
Table 16 . Experimental and calculated parameters of pseudo-first and second order kinetic models of Pb2+ (pH=4) adsorbed on CA and CACs powder.
qe Experimental Pseudo first order Pseudo second order
qe calculated
K1
R2 qe
calculated
K2
R2
CA 43.45 9.698 0.1718 0.9431 44.09 4.7x10-6 0.9997
CAB
34.55
13.06
0.1026
0.9001
35.95
3.63x10-5
0.9967
CAMK
38.92
27.24
0.1758
0.8219
40.7
2.34x10-5
0.9978
CAHA(I)
43.39
32.13
0.1921
0.9459
45.74
1.53x10-5
0.9955
CAHA(II) 42.45 15.99 0.1938 0.9859 44.3 1.47x10-5 0.9974
CAD
32.21
14.73
0.107
0.9513
34.16
5.62x10-5
0.9944
CAG
44.1
21.03
0.1394
0.6528
44.62
1.18x10-5
0.9932
Table 17: Experimental and calculated parameters of pseudo-first and second order kinetic models of Fe3+ (natural pH of 2.6) adsorbed on CA and CACs powder.
qe experimental Pseudo first order Pseudo second order
qe calculated
K1
R2 qe
calculated
K2
R2
CA
47.09
43.46
0.2330
0.8915
49.26
9.57x10-6
0.9995
CAB
44.75
39.40
0.2171
0.9256
47.12
1.37x10-6
0.9963
CAMK
44.95
29.695
0.2163
0.8390
46.25
6.56x10-6
0.9997
CAHA(I)
49.72
1.48
0.1037
0.8751
50
2.2x10-3
0.9999
CAHA(II)
49.39
1.41
0.0747
0.7308
50
2.2x10-3
0.9999
CAD
43.4
21.32
0.1857
0.8184
44.3
9.63x10-6
0.9974
CAG
45.75
5.47
0.0938
0.6573
46.08
3.18x10-6
0.9981
Figure 44: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4)and Fe3+ (pH=2.6) adsorbed on CA at experimental conditions of 100 mg/L metal ion concentration, CA mass 0.02 g/10 mL and 25±1oC.
Figure 45: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAB at experimental conditions of 100 mg/L metal ion concentration, CAB mass 0.02 g/10 mL and 25oC.
Figure 46: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAMK at experimental conditions of 100 mg/L metal ion concentration, CAMK mass 0.02 g/10 mL and 25±1oC.
Figure 47: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAHA(I) at experimental conditions of 100 mg/L metal ion concentration, CAHA(I) mass 0.02 g/10 mL and 25±1oC.
Figure 48: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAHA(II) at experimental conditions of 100 mg/L metal ion concentration, CAHA(II) mass 0.02 g/10 mL and 25±1oC.
Figure 49: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAD at experimental conditions of 100 mg/L metal ion concentration, CAD mass 0.02 g/10 mL and 25±1oC.
Figure 50: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAG at experimental conditions of 100 mg/L metal ion concentration, CAG mass 0.02 g/10 mL and 25±1oC.
from all the obtained results illustrated in Figures 37-50 and Tables 15-17, it is obvious that the regression coefficient (R2) from pseudo-second order rate equation for all the adsorbents is higher than that of the pseudo-first order model. On the basis of the regression coefficient and calculated values of adsorption capacity, the adsorption process is found to obey and exhibited best fit to the pseudo-second-order kinetic model which mean that the rate-limiting step might be chemical adsorption or chemisorption involving valency forces through exchange of electrons between the sorbate and the sorbent, also only one ion of the metal is sorbed onto two sorption sites on the sorbent surface [192, 193].
3.4.3. Prediction of adsorption rate-limiting step
There are essentially three consecutive mass transport steps associated with the adsorption of solute from the solution by an adsorbent. These are (1) film diffusion, (2) intraparticle or pore diffusion, and (3) sorption into interior sites. The third step is very rapid and hence, film and pore transports are the major steps controlling the rate of adsorption [179, 194].
The most commonly used technique for identifying the mechanism involved in the adsorption process is by fitting an intraparticle diffusion plot [195]. The amount of metal ions adsorbed (qt) at time (t), is plotted against the square root of t (t1/2), according to Eq. proposed by Weber and Morris as follows:
Qt = Kid t0.5 + C
where C is constant and kid is the intraparticle diffusion rate constant (mg/g min1/2), qt is the amount adsorbed at a time (mg/g), t is the time (min), and kid (mg/g min1/2) is the rate constant of intraparticle diffusion. Due to the varying extent of adsorption in the initial and final stages of the experiment two straight lines with different slopes are obtained (Figures 51-57).
The two regions in the qt vs. t0.5 plot suggest that the sorption process proceeds by surface sorption and intraparticle diffusion. The initial rapid uptake can be attributed to the boundary layer effects (film diffusion). After the external surface loading was completed, the intraparticle diffusion or pore diffusion takes place, The second linear part of the plot presented in Fig.51-57, corresponds to the transportation of Pb2+ and Fe3+ within CA and CACs particles [196]. The slope of the second linear portion of the plot has been defined to yield the intraparticle diffusion parameter of ki1, Ri12 (first stage) and ki2, Ri22 (second stage) are listed in Tables (18-20). On the other hand, the intercept of the plot give an idea about the thikness of boundary layer effect [197]. The larger the intercept, the greater the contribution of the surface sorption in the rate-controlling step [195].
As it can be seen in Figures (51-57), the plot indicated that the intraparticle diffusion was not the rate-controlling step because it did not pass through the origin [196]. The deviation of the straight lines from the origin may be due to the difference in the rate of mass transfer in the initial and final stages of adsorption [198]. Further, the first straight portion is attributed to a macropores diffusion process and the second linear portion can be ascribed to a micropore diffusion process [192, 199]. In addition, it is clear from Fig. (51-57) that the first stage is faster than the second one. This behaviour may be correlated with the very slow diffusion of the adsorbate from the surface film into the micropores, which are the least accessible sites for adsorption [197].
Figure 51: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+(pH=2.6) on CA.
Figure 52: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) on CAB composite.
Figure 53: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) on CAMK composite.
Figure 54: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+( pH 4) and Fe3+ (pH=2.6) on CAHA(I) composite.
Figure 55: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) on CAHA(II) composite.
Figure 56: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) on CAD composite.
Figure 57: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) on CAG composite.
Table 18: The parameters of intraparticle diffusion model for adsorption of Pb2+ (pH=5.7) onto CA and CACs.
Intrapartical diffusion modle
Ki1
(mg/g m0.5)
Ki2
(mg/g m0.5)
Ri12
Ri22
Intercept i1
Intercept i2
CA
0.617 0.071 0.9541 0.7161 37.55 40.50
CAB
1.460 0.13 0.8094 0.7094 24.3 32.07
CAMK
1.682 0.049
0.9365 0.6601 25.94 34.31
CAHA(I)
3.53 0.159 0.9989 0.6689 23.43 38.85
CAHA(II)
2.087 0.019 0.9917 0.8957 25.62 36.93
CAD
3.77 0.15 0.7123 59.22 15.48 33.45
CAG
1.36 0.12 0.9615 0.8696 33.92 41.05
Table 19: The parameters of intraparticle diffusion model for adsorption of Pb2+ (pH=4) onto CA and CACs.
Intrapartical diffusion modle
Ki1
(mg/g m0.5)
Ki2
(mg/g m0.5)
Ri12
Ri22
Intercept i1
Intercept i2
CA
1.11 0.066 0.9056 0.4133 38.07 42.87
CAB
1.71 0.10 0.9716 0.7251 33.6 24.54
CAMK
1.88 0.021 0.9745 0.9251 28.48 38.74
CAHA(I)
3.21 0.086 0.8603 0.3367 27.94 42.06
CAHA(II)
2.01 0.039 0.9556 0.9346 31.56 42.12
CAD
2.25 0.14 0.9119 0.7920 19.74 30.89
CAG
1.43 0.025 0.7045 0.4308 34.97 43.84
Table 20: The parameters of intraparticle diffusion model for adsorption of Fe3+ (pH=2.6) onto CA and CACs.
Intrapartical diffusion modle
Ki1
(mg/g m0.5)
Ki2
(mg/g m0.5)
Ri12
Ri22
Intercept i1
Intercept i2
CA
2.75 0.097 0.9892 0.108 33.98 46.16
CAB
3.79 0.11 0.7902 0.1968 27.28 43.7
CAMK
1.63 0.012 0.9593 0.9589 36.36 44.85
CAHA(I)
0.3 0.0017 0.6592 0.3204 48.14 49.65
CAHA(II)
0.17 0.013 0.9673 0.5278 48.37 49.25
CAD
1.33 0.020 0.9937 0.9134 35.9 43.24
CAG
0.62 0.062 0.9351 0.7898 41.65 45.22
from Tables (18-20) it can be conclouded that for adsorption of Pb2+(pH=5.7), Pb2+ (pH=4) and Fe3+ (pH=2.6) onto CA and CACs, the calculated values of ki1 were higher than that of ki2. The reason could be as circumscription of the available vacant space for diffusion in them, so of pore blockage. The values of the correlation regression coefficients characterizing the applicability of the intraparticle diffusion model (Ri12, Ri22) were lower than that of R2 (pseudo second order), but commensurable with R12 ( pseudo first order). Actually, the three models stated above could describe the proposed sorption process to a definite extend, but they could not predict the high rate of adsorption during the first minutes of the process. Probably, the initial stages are controlled by external mass transfer or surface diffusion, followed by chemical reaction or a constant-rate stage, and diffusion causing gradual decrease of the process rate [192]
3.5. Adsorption isotherms:
Adsorption isotherm studies are necessary for illustrating the adsorption process at equilibrium conditions. An adsorption isotherm is characterized by certain constants which express the adsorbent affinity and can also be used for finding the adsorption capacity of the sorbent. The adsorption of iron and lead from polluted water using CA and CACs could be assumed to have abehavior fitting with the isothermal adsorption model in which the adsorbate keeps a dynamic equilibrium between the adsorotion and desorption at afixed temperature [178, 200].
Two most widely used mathematical models Langmuir and Freundlich adsorption isotherms are adopted for expressing the quantitative relationship between the extent of sorption and the residual solute concentration. Langmuir adsorption isotherm assumes monolayer coverage of adsorabate over ahomogeneous adsorbent surface and the adsorption of each molecule onto the surface has the same activation energy of adsorption.
Freundlich adsorption isotherms assumes aheterogeneous surface with anon-uniform distribution of heat of adsorption over the surface with the possibility of the multilayer adsorption [178, 179]. The maximum metal ions adsorption capacities are determined by analyzing the experimental data for heavy metal adsorption onto CA and CAHA(I) [201], as they provide the higher removal effeciency. The data of Pb2+ and Fe3+ adsorption by CA and CAHA(I) are examined in accordance with langmiur adsorption isotherm models whose linearized equation was:
Ce/qe =Ce/qm + 1/ KL qm
where qm (mg /g) and KL (L /mg) are constants in Langmuir’s equation which are referred to the maximum adsorption capacity corresponding to complete monolayer coverage and the Langmuir model constant that is indirectly related to the energy of adsorption. Also qe and Ce parameters represent the equilibrium adsorption capacity and the equilibrium concentration heavy metal ion that is remaining in solution, respectively. qe is calculated as follows:
qe =((Co-Ce)V)/m
where Co is the initial metal ion concentration (mg/L), Ct is the equilibrium concentration of adsorbate (mg/L) (mg/L), V is the initial solution volume (L) and m is the adsorbate dose (g). A plot of Ce/qe vs. Ce (Fig. 58-61) gives a linear trace with a slope of 1/qm and intercept of 1/ KL qm. A further analysis of the Langmuir equation can be made on the basis of a dimensionless equilibrium parameter, RL, also known as the separation factor,
given by
RL=1\(1+ KLCe)
where Ce is equilibrium liquid phase concentration of the solute at which adsorption is carried out. The value of RL lies between 0 and 1 for favorable adsorption, while RL > 1 represents unfavorable adsorption, and RL = 1 represents linear adsorption, while the adsorption process is irreversible if RL = 0 [179, 202].
Also the obtained data are examined in accordance with the Freundlich equation which is expressed as
qe =log KFCe1/n
where KF is the measure of adsorption capacity and n is the adsorption intensity linear form of Freundlich
log qe =log KF + (1/n) log Ce
where qe is the amount adsorbed (mg/g), Ce is the equilibrium concentration of adsorbate (mg/L), and KF and n are the Freundlich constants related to the adsorption capacity and adsorption intensity, respectively. A plot of log qe vs. log Ce (Fig. 58-61) gives a linear trace with a slope of 1/n and intercept of log KF. The 1/n value in the range of 0 and 1 is a predicting of adsorption intensity of metal ion onto the adsorbent and the type of isotherm to be irreversible (1/n=0), favourable (0<1/n<1) and unfavourable (1/n >1) [197], the 1/n value determine the surface heterogeneity, becoming more heterogeneous as its value gets closer to zero. In addition, the value of n varies with the heterogeneity of the adsorbent, if n < 10 and n > 1 indicating the adsorption process is favorable [193, 203].
The isotherm parameters and correlation coefficients calculated for the adsorption of Pb2+ and Fe3+ using CA and CAHA(I) are listed in Tables (21,22).
Table 21: Isotherm parameters and correlation coefficients calculated for the adsorption of Pb2+ using CA and CAHA(I).
Adsorbent
Langmuir Isotherm
Freundlich Isotherm
qmax(mg/g) KL R2 KF 1/n N R2
CA
51.78
1.264
0.9353
18.62
0.377
2.64
0.6028
CAHA(I)
52.99
0.702
0.8652
16.89
0.3450
2.89
0.9555
Table 22: Isotherm parameters and correlation coefficients calculated for the adsorption of Fe3+ using CA and CAHA(I).
Adsorbent
Langmuir Isotherm
Freundlich Isotherm
qmax(mg/g) KL R2 KF 1/n N R2
CA 66.53 0.0816 0.9361 15.84 0.2602 3.84 0.9656
CAHA(I) 113.63 0.1560 0.7675 24.35 0.3334 2.99 0.9178
from Table 21, it can be concluded that for adsorption of lead on CA, the Langmuir isotherm (R2 > 0.93) fitted the experimental results better than those of the Freundlich isotherm (R2 > 0.60) as reflected with the correlation coefficient, indicating the homogenous feature presented on the CA surface and demonstrates the formation of monolayer coverage of the lead ions on the CA surface, the adsorption is localized, all active sites of surface have similar energies and no interaction between adsorped molecules. Moreover, the value of RL was 0.0003. This also suggests an irreversible adsorption between CA and Pb2+ ions [202].
On the other hand, for CAHA, it can be stated that the Freundlich isotherm (R2 > 0.95) fitted the experimental results comparable to the Langmuir isotherm (R2 > 0.86), indicating that the adsorbed amount increased with initial concentration. The slope 1/n provides information about surface heterogeneity and surface affinity for the solute. As a higher value of 1/n (0.34) is obtained, it corresponds to the greater heterogeneity of the adsorbent surface. Furthermore, the value of n > 1 obtained from the Freundlich isotherm indicating (2.8), that this process is also favorable [203] and heterogeneous sorption. The maximum adsorption capacities of the Pb2+ ions are found to be 51.78 and 52.99 mg/g for CA and CAHA(I), respectively.
from Table 22, it can be stated that for adsorption of iron on CA and CAHA(I), the Freundlich isotherm (R2 > 0.96), (R2 > 0.91) fitted the experimental results comparable to the Langmuir isotherm (R2 > 0.93), (R2 > 0.76) for CA and CAHA(I) respectively. The slope 1/n provides information about surface heterogeneity and surface affinity for the solute. As a higher value of 1/n is obtained, it corresponds to the greater heterogeneity of the adsorbent surface. Furthermore, the value of n > 1 obtained from the Freundlich isotherm (3.8, 2.9), indicating that this process is also favorable and heterogeneous sorption.
Firure 58: Langmuir and Freundlish isotherms for adsorption of Pb2+ on CA Powder.
Firure 59: Langmuir and Freundlish isotherms for adsorption of Pb2+ on CAHA(I) Powder.
Firure 60: Langmuir and Freundlish isotherms for adsorption of Fe3+ on CA Powder.
.
Firure 61: Langmuir and Freundlish isotherms for adsorption of Fe3+ on CAHA(I) Powder
SUMMARY AND CONCLUSION
The present thesis comprises of there chapters:
Chapter (1) includes introduction and literature review that focused on the heavy metals found in industrial effluents as hazardous pollutants that may affect hardly the environment.
This chapter also includes various technologies which have been used to remove metal ions. Especial attention is given to two types of heavy metals ions (Pb2+ and Fe3+), their abundance, use in several industries and harmful effect on the environment especially to aquatic systems.
The literature review also focused on calcium alginate and its biomedical applications, also focused on bentonite, metakaolin, diatomeous silica, gelatin and hydroxy apatite (produced from biowastes) which form composites with calcium alginate and their role in heavy metals removal.
Chapter (2) includes experimental and methods which focused on types of chemical used, methods of preparation of calcium alginate, hydroxy apatite and calcium alginate composites and characterization tecniques such as X-Ray Diffraction (XRD), Fourier Transformer Infrared (FTIR), Scanning Electron Microscope(SEM) and Energy Dispersive Analysis X-ray (EDX). Preparation of heavy metal solutions with the assessment of their concentrations before and after adsorption process using Perkin Elmer 2380 Atomic Absorption Spectro Photometer. The six synthesized composite beside calcium alginate used in the present study are listed in the following table:
Composite Compound Source of raw material Abbrev.
1 Calcium alginate Oxford Lab. Reagent CA
2 Calcium alginate-Bentonite Bentonite (Abu Zaabal Fertilizer & Chemicals Co. CAB
3 Calcium alginate-Metakaolin Metakaolin (Sinai Peninsula) CAMK
4 Calcium alginate-HAP(1) HAP (egg shell calcined at 900oC) CAHA(I)
5 Calcium alginate-HAP(2) HAP (bovin bone calcined at 1000oC) CAHA(II)
6 Calcium alginate-Diatomeous Diatomeous (Kazakhstan) CAD
7 Calcium alginate-Gelatin Oxford Lab. Reagent CAG
Chapter (3) is concerned with the results and discussion that includes characterization of the prepared composites before and after heavy metal uptake using XRD, FTIR, SEM and EDX techniques. It also include the adsorption process of heavy metal uptake and kinetics studies of the adsorption process.
I. Fourier Transform Infrared (FTIR) analysis:
The FTIR spectra of calcium alginate showed important absorption bands regarding hydroxyl, ether and carboxylic functional groups. Stretching vibrations of O–H bonds of alginate appeared at 3444 cm-1. In the spectra of CAG, the absorption band at around 3442 cm-1 concerned with OH stretching vibration for CA slightly broadened and shifted to a lower wave number with the blending with gelatin, suggesting the formation of an intermolecular hydrogen bond.
CA raw material, reveals asimilar spectra with CAB. A strong and sharp band is detected at 1022 cm-1 which is related to Si–OH stretching vibrations. 1034 cm-1 assigned to Si-O stretching, and 875 cm-1 is due to OH bending of the Al-Al-OH group. A similar OH bending vibration is observed for Al-Mg-OH at 842 cm-1, 690 cm-1 assigned to quartz. Also, there is a shoulder peak at 520 cm-1 (Al-O-Si bending), and 464 cm-1 (Si-O-Si bending). The band of Si-O-Si bending in CAD is the strongest and sharpest compared to CAMK and CAB.
The bands corresponding to the samples of CAHA(I) and CAHA(II) which appears at 3570 and 630-633 cm-1 (due to structural OH of hydroxy apatite) confirm the hydroxyapatite. The most intensive bands in the range of 1044 –1090 cm-1 corresponded to the triply degenerated asymmetric stretching vibrations of P-O. Otherwise the peak at 962.97 cm-1 indicates the non-degenerated asymmetric mode of PO43-. The very strong and sharp bands observed at 569-572 and 602-603 cm-1 attributed to triply degenerated bending mode of the O-P-O in PO43- group. The larges parting distance of these bands revels the crystalline phase.
II. X-ray diffraction for crystal phase detection:
The X-ray diffraction pattern of CA and CACs in the range of 2θ = 5-60o diffraction degree shows two typical peaks in 2θ =16° and 22° corresponding to calcium alginate. The XRD of CAG shows typical peaks around 12° and 21°. The samples of CAB, CAMK and CAD revealed very similar XRD pattern peaks of quartz phase.
The sample CAHA(I) and CAHA(II) revealed very similar XRD pattern peaks of 100 % HAP (reference code :01-086-1194) in the rang of 2θ =31o and 32o and d (Ao) value spacing equal 2.81 and 2.78.
XRD analysis of CAHA(II) also indicated the absence of secondary phases, such as tri calcium phosphate (TCP) or calcium oxide (CaO). In the case of CAHA(I), its diffraction pattern revealed additional phase of β-tricalcium phosphate, beside the HAP as a main phase. Hydoxyapatite prepared from bovine bone contain certain amount of carbonate (CO32-) in its lattice structure. The carbonate ions affect on the the degree of crystallinite and hence increase the bioactivity of pure HAP.
III. Scanning Electron Microscope (SEM) and Energy Dispersive X-ray (EDX)
According to the results obtained by EDX analysis, carbon, oxygen and calcium are the major constituents of all the prepared composites related to calcium alginate polymer present. EDX analyses of CA and CAG revealed that the major elements present are carbon and oxygen corresponding to polymer composition. The SEM of CAG showed a smooth and homogeneous morphology, suggesting high miscibility and blend homogeneity between calcium alginate and gelatin.
In the case of CAB, CAMK and CAD, the presence of an obvious peak related to the Si compounds is evident. According to the results obtained by EDX analysis, weigh percent of elements present indicating a large portion of the composites is composed of Si compounds which is suitable for an efficient sequestering metal cations from aqueous solution.
EDX analyses of the prepared HAP revealed that inorganic phases of bovin bone and egg shells were mainly composed of calcium and phosphorus as the major constituents with some minor components such as C, O, Na, Mg and Si. The weight and atomic percentage shows that the Ca/P ratio around 1.7 and 1.8 which is below 2 and acceptable where the ideal Ca/P ratio of HA is 1.67, in HAP composites (CAHA(I) and CAHA(II)) this ratio increased than 2 due to the excess amount of Ca crosslinkage of calcim alginate presents.
IV. FTIR Analysis after metal ion uptake:
FTIR spectra of CA and CAHA(I) after targeted metal ions uptake ( Pb2+, Fe3+ ) showed that there is no significant change in peak positions after metal ion uptake and also no new absorption peaks are detected . There are alittle peak shifts which may be attributed to the corporation and substitution of metal ion in the lattice structure of CA and CAHA(I).
V. X-ray diffraction for crystal phase detection after metal ion uptake:
XRD patterns of CA and CAHA(I) after Pb2+ and Fe3+ metal ions removal did not revealed any new phases. supported the proposal that Pb2+ and Fe3+ ions uptake was not dependent on dissolution/precipitation mechanisms. Pb2+ and Fe3+ ions removal may be occurs by adsorption mechanisms like surface complexation or ionic exchange. The XRD patterns showed some changes in their relative intensities and crystal sizes. Also, ther are some little shifts in d- spacing values, this may be due to the ion exchange between Ca2+ and metal cations of Pb2+ and Fe3+ in lattice structure of CA and CAHA(I) Powder.
VI. Scanning Electron Microscope (SEM) and Energy Dispersive X-ray (EDX) after metal ion uptake:
SEM of CA and CAHA(I) after metal ion uptake revealed some changes in morphology and microstructure. Morever EDX results indicated the presence of Pb2+ and Fe3+ ions with CA and CAHA(I). the peak of Fe3+ is more intense than that of Pb2+ and this confirms, the higher value of Fe3+ ions uptake by CA and CAHA(I) comparing with Pb2+ ions.
The weight and atomic percentage of Ca2+ ions in CA and CAHA(I) after metal ion uptake was less than its value before metal ion removal. The decrease in calcium percentage may be attributed to ion exchange between targeted metal ions and; (1)calcium ions of calcium alginate in CA and CAHA(I), (2) or Ca ions of HA present in CAHA(I). This ion exchange mechanism between Pb2+ ions (as example) and Ca2+ ions of HA produced anew phase of hydroxypyromorphite. However , this phase is not detected in the present study this may be attributed to under limit of XRD.
Also the uptake process may be occurs during chelation bonding of targeted metal ions with two carboxylic groups of alginate and one or two OH sites of the alginate. In this case metal ions may forms complexes with two adjacent alginate rings. Here,‘‘adjacent’’ means either two neighbor alginate rings of a single polymeric chain (intramolecular chelation) or two rings from two parallel chains (intermolecular chelation).
VII. Metal ion uptake by adsorption process.
1- Effect of contact time on the adsorption process.
The effect of contact time on the adsorption capacity of CA and CACs for Pb2+ ( natural pH of 5.7), Pb2+ (pH=4) and Fe3+ ( natural pH of 2.6) indicated that, the equilibrium points of adsorption are attained within the first 10 – 60 min. of contact time with different removal efficiency and slightly similar exponential phase as listed in table below. the adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+(pH=2.6)onto CA and CACs powder was rapid in the initial stages and was almost same at high contact time. This behavior could be attributed to the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption process. Further increase in contact time led to no significant adsorption of metal ions by the adsorbents probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. As it can be seen in the following table, CAG is the higher efficiency for lead (Pb2+) removal then CAHA(I) and CA. On the other hand, CAHA(I) is the higher efficiency for iron (Fe3+) removal then CAHA(II) and CA.
Equilibrium point time (min.) Removal effeciency(%) Exponential phase time (min.)
Pb2+
(pH=5.7)
Pb2+
(pH=4)
Fe3+
(pH=2.6)
Pb2+
(pH=5.7)
Pb2+
(pH =4)
Fe3+
(pH=2.6)
Pb2+
(pH=5.7)
Pb2+
(pH=4)
Fe3+
(pH=2.6)
CA
60
30
30
82.3
86.82
94.15
5-60
5-30
5-30
CAB
60
30
30
66.4
68.2
89.42
5-60
5-30
5-30
CAMK
30
30
30
69.26
77.7
89.87
5-30
5-30
5-30
CAHA(I)
60
30
30
80.78
86.62
99.33
5-60
5-30
5-30
CAHA(II)
30
30
10
74.08
84.67
98.65
5-3
5-30
5-10
CAD
30
30
30
68.96
63.23
86.72
5-30
5-30
5-30
CAG
30
30
30
83.43
87.93
91.11
5-30
5-30
5-30
2- Effect of pH
The effect of pH on the adsorption is performed only for Pb2+ because of Fe3+ solution was stable only at pH lower than 3. The study is achieved with two pHs values (4 and 5.7) the original solution of lead is at pH=5.7 and pH=4 at 5-120 minutes contact time, 20 mg dosage of the different composites and 100 ppm of metal solution at 25±1°C. The adsorption efficiency of Pb2+ at pH=4 is higher than that of pH 5.7 for all the composites.
3- Effect of adsorbent dosage on metal ion adsorption.
The experimental results of the adsorption of Pb2+ on CA ( as astandard model ) as afunction of adsorbent dosage 10, 15 and 20 mg/10 mL, initial Pb2+ concentration of 100 mgL-1, natural pH of 5.7, temperature 25oC at the optimal contact time (30 min) and interval contact time ( 5-30 min) showed that, the Pb2+ adsorption percent rapidly increased with the increase in the adsorbent dosage . this can be attributed to higher adsorbent dosage due to the increased surface area providing more adsorption sites available which gave rise to higher removal of lead.
VIII. kinetics studies of the adsorption process.
The kinetic study is useful to predict the adsorption rate which is very important in modeling and designing of the adsorption process. The kinetic of adsorption are evaluated at an initial concentration of 100 mg/L for Pb2+(pH=5.7),Pb2+( pH 4)and Fe3+(pH=2.6), adsorbent dosage of 0.02 g/10 mL and temperature of 25oC. By appling the pseudo-first rate equation of lagergren, it is clear that the regression coefficient does not close to unity. Also, the values of qe obtained from pseudo-first order equation for all the adsorbent are different and not matched notably with the experimental qe value. from the linear plots of pseudo-second rate equation of lagergren, the qe,experimental and the qe,calculated values are very close to each other, and also, the calculated coefficients of determination, R2, are close to unity
from all the obtained results, it is obvious that the regression coefficient (R2) from pseudo-second order rate equation for all the adsorbents was higher than that of the pseudo-first order model. On the basis of the regression coefficient and calculated values of adsorption capacity, the adsorption process was found to obey and exhibited best fit to the pseudo-second-order kinetic model which is mean that the rate-limiting step might be chemical adsorption or chemisorption involving valency forces through exchange of electrons between the sorbate and the sorbent, also only one ion of the metal is sorbed onto two sorption sites on the sorbent surface.
IX. Prediction of adsorption rate-limiting step
There are essentially three consecutive mass transport steps associated with the adsorption of solute from the solution by an ads0rbent. These are (1) film diffusi0n, (2) intraparticle or p0re diffusion, and (3) sorption into interior sites. The third step is very rapid and hence, film and pore transports are the major steps controlling the rate of adsorption. The most commonly used technique for identifying the mechanism involved in the adsorption process is by fitting an intraparticle diffusion plot proposed by Weber and Morris.
The results stated that the sorption process proceeds by surface sorption and intraparticle diffusion. The initial rapid uptake can be attributed to the boundary layer effects (film diffusion). After the external surface loading was completed, the intraparticle diffusion or pore diffusion takes place. However, the plot indicated that the intraparticle diffusion was not the rate-controlling step because it did not pass through the origin.
X. Adsorption isotherms:
Adsorption isotherm studies are necessary for illustrating the adsorption process at equilibrium conditions. Two most widely used mathematical models Langmuir and Freundlich adsorption. Langmuir adsorption isotherm assumes monolayer coverage of adsorabate over ahomogeneous adsorbent surface and the adsorption of each molecule onto the surface has the same activation energy of adsorption. Freundlich adsorption isotherms assumes aheterogeneous surface with anon-uniform distribution of heat of adsorption over the surface with the possibility of the multilayer adsorption
The results of single metal ion adsorption of Pb2+ onto CA at 25 oC can be represented well by langmiur than Freundlish model with good correlation coefficient (R2). This means that the adsorbates containing Pb2+ was adsorbed in such amanner that only one atomic layer of adsorbate can be adsorbed and distributed uniformly on the surface of the adsorbents (CA) and the adsorption of each molecule onto the surface has the same activation energy of adsorption. the value of RL was 0.0003. This also suggests an irreversible adsorption between CA and Pb2+ ions.
For iron and lead adsorbed onto CAHA and iron adsorbed into CA, it can be stated that the Freundlich isotherm well fitted the experimental results comparable to the Langmuir isotherm indicating that the adsorbed amount increased with initial concentration. The slope 1/n provides information about surface heterogeneity and surface affinity for the solute. As a higher value of 1/n is obtained, it corresponds to the greater heterogeneity of the adsorbent surface. Furthermore, the value of 1 < 1/n > 0 and the value of n > 1 obtained from the Freundlich isotherm indicating, that this process is also favorable and heterogeneous sorption.
from all the obtained results and analysis we can stated that, the uptake of Pb2+ and Fe3+ may be occurs by adsorption mechanisms like surface complexation during chelation bonding of targeted metal ions with two carboxylic groups of alginate and one or two OH sites of the alginate ring forming complexes with two adjacent alginate rings. Here,‘‘adjacent’’ means either two neighbor alginate rings of a single polymeric chain (intramolecular chelation) or two rings from two parallel chains (intermolecular chelation). or ion exchange between targeted metal ions and; (1) Calcium ions of calcium alginate in CA and CACs, (2) or Ca ions of HA present in CAHA(I) and CAHA (II).
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الملخص العربي
الرساله تنقسم الي ثلاثة فصول رئيسية :
الفصل الاول : ويتضمن المقدمة والمراجع التاريخية التي تخص العناصر الثقيلة كمُلوثات بيئية والتي تنتج بكمية كبيرة من المخلفات الصناعية والتي تؤثر بصورة خطيرة علي البيئة وصحة الانسان. وشمل هذا الفصل أيضا مختلف التقنيات المستخدمه لأزالة هذه العناصر. وتم القاء الضوء علي عنصرين من هذه العناصر وهما الرصاص والحديد من حيث مصادر تواجدها من المخلفات الصناعية وتأثيرها الضارعلي البيئة المائية. ويتضمن هذا الفصل أيضا شرح وافي لألجينات الكالسيوم وتطبيقاتها في ازالة العناصر الثقيلة, والجيلاتين وبعض الطفلات مثل البنتونيت والميتاكاولين وسيليكا الدايتموس والهيدروكسي اباتيت المحضرة من المخلفات الحيوية (مثل قشر البيض وعظام البقر) والتي تكون مركبات مع الجينات الكالسيوم للتخلص من العناصر الثقيلة.
الفصل الثاني : ويشمل الكيماويات المستخدمة , وطرق تحضير ألجينات الكالسيوم والهيدروكسي اباتيت ومركبات ألجينات الكالسيوم المختلفة وطرق التعرف عليها باستخدام قياسات حيود الاشعة السينية , الأشعة تحت الحمراء, الميكرسكوب الألكتروني الماسح للضوء وطاقة الاشعة السينية المشتتة. وكذلك طرق تحضير محاليل ايونات العناصر الثقيلة وتقديرتركيزها قبل وبعد عملية الإمتزاز باستخدام جهاز الامتصاص الذري. ويتضمن الجدول التالي مصادر المخلفات الحيوية والمواد الخام وألجينات الكالسيوم ومركباته السته المحضرة :
Composite Compound Source of raw material Abbrev.
1 Calcium alginate Oxford Lab. Reagent CA
2 Calcium alginate-Bentonite Bentonite (Abu Zaabal Fertilizer & Chemicals Co. CAB
3 Calcium alginate-Metakaolin Metakaolin (kaolin (Sinai Peninsula) calcined at 800oC CAMK
4 Calcium alginate-HAP(1) HAP (egg shell calcined at 900oC) CAHA(I)
5 Calcium alginate-HAP(2) HAP (bovin bone calcined at 1000oC) CAHA(II)
6 Calcium alginate-Diatomeous Diatomeous (Kazakhstan) CAD
7 Calcium alginate-Gelatin Oxford Lab. Reagent CAG
الفصل الثالث : ويتناول النتائج التي تم الحصول عليها ومناقشتها من حيث توصيفها قبل وبعد عملية الامتزاز باستخدام اجهزة SEM وEDX مع XRD و FTIR ويشمل ايضا دراسة عملية الامتزاز و حركية عملية الامتزاز.
التعرف علي المجموعات الوظيفية من قياسات الاشعة تحت الحمراء(i)
اظهرت تحاليل عينة الجينات الكالسيوم تواجد حزم امتصاص هامة تعود الي المجموعات الوظيفيه الخاصة ب الهيدروكسيل والكربوكسيليك والايثر. حيث تظهر مجموعة الهيدروكسيل الخاصه ألجينات الكالسيوم CAعند طول موجي حوالي 3444 cm-1 , ويظهر طيف مركب الجينات الكالسيوم مع الجيلاتين CAG أن ذروة الامتصاص عند حواليcm-1 3442 والخاصة باهتزاز مجموعة الهيدروكسيل OH لـ ألجينات الكالسيوم قد اتسعت قليلاً وتحركت إلى طول موجي أقل بالمزج مع الجيلاتين ، مما يشير إلى تكوين رابطة هيدروجينية بين جزيئية.
يوجد تشابه كبير بين طيف ألجينات الكالسيوم وطيف ألجينات الكالسيوم مع البنتونيت CAB والميتاكاولين CAMK وسيليكا الدايتاموس CAD , ففي CAB تظهر ذروة امتصاص حادة وقوية عندcm-1 1022 والخاصة باهتزاز مجموعة Si-OHوعند cm-1 1034 لمجموعة Si-O وعند 875 cm-1 بسبب وجود OH bending لمجموعةAl-Al-OH كذلك تهتز مجموعة مشابهه منOH bending , خاصة ب Al-Mg-OH تظهر عند842 cm-1 و cm-1 690 تعود الي تواجد الكوارتز. ايضا يوجد حزمة امتصاص شولدر عند 520 cm-1 تعود الي (Al-O-Si bending ) وعند 464 cm-1 خاصة ب (Si-O-Si bending). تلاحظ ايضا ان حزمة الامتصاص الخاصة ب Si-O-Si bending في عينة مركب ألجينات الكالسيوم مع الداياتاموس تكون اقوي وأحد مقارنة بعينات CAB و CAMK.
بالنسبة لعينات ألجينات الكالسيوم مع الهيدروكسي أباتيت المحضرة من قشور البيض وعظام البقر CAHA(I) و CAHA(II) فتظهر حزم امتصاص عند 3570 cm-1 و 630-633 cm-1 بسبب اهتزازمجموعة OH للهيدروكسي أباتيت. كما أظهرت النتائج ايضا أن ذروة الامتصاص الاكثر شده في المدي من 1044 cm-1 الي 1090 cm-1 ترجع الي اهتزازات الرابطة P-O في مجموعة الفوسفات وتكون تللك الاهتزازات متماثلة , اما عند 962.97 cm-1 فتكون تلك الاهتزازات لمجموعة الفوسفات غير متماثلة. حزم الامتصاص القوية جدا والحادة والتي تظهر عند 569-572 cm-1 و عند 602-603 cm-1 فتعزي الي اهتزازات مجموعات O-P-O في مجموعة الفوسفات PO43-.
(ii) قياسات حيود الاشعه السينية :
بينت التحاليل باستخدام X-ray diffraction لعينات ألجينات الكالسيوم ومركباته في في المدي 5-60o=θ2 تواجد عدد ذروتين خاصة بالجينات الكالسيوم في نطاق =16 o , 22 oθ2 واظهرت النتائج ايضا تواجد الجيلاتين في عينة CAG في نطاق =12 o , 21 oθ2 وتظهر نتائج تحاليل عينات CAB, CAD, CAMK تشابه كبيربينهم و تكون الطور الخاص بالكوارتز.
كما تظهر النتائج تشابه كبير لعينات الجينات الكالسيوم مع الهيدروكسي اباتيت المحضر من قشور البيض CAHA(I) و CAHA(II) ذات الهيدروكسي اباتيت المحضر من عظام الابقار وتكون طور الهيدروكسي اباتيت طبقا للكود المرجعي 01-086-1194 في نطاق =31 o , 32 oθ2 عند قيم مسافات تباعد (d-spacing) مساوية 2.81 و 2.78انجستروم , ولا تحتوى عينات CAHA(II) على أي أطوار اخرى مثل أكسيد الكالسيوم CaO أوفوسفات الكالسيوم في حين ان عينات CAHA(I) تظهر تواجد طور فوسفات الكالسيوم بجانب طور الهيدروكسي اباتيت كطور رئيسي وتظهر نتائج تحاليل عينة الجينات الكالسيوم مع الهيدروكسي اباتيت المحضرة من عظام الابقار CAHA(II) انها تحتوي علي كمية من الكربونات في الشبكة البلورية. وتوفر ايونات الكربونات زيادة في النشاطية الحيوية للهيدروكسي أباتيت كما تؤثر في درجة التبلور للمركب.
(iii) التعرف على الشكل المورفولوجى والتحليل النوعي والكمى للعينات المحضرة بأستخدام قياسات الميكروسكوب الإلكترونى الماسح للضوء وطاقة الأشعة السينية المشتتة.
أظهرت صور الميكروسكوب الإلكتروني لجميع العينات تكتل بلوري للجزيئات مع شكل غير منتظم
نسبيا ، مع أحجام مختلفة من البللورات وتوضح الصور الشكل المورفولوجي المتجانس والناعم لعينة الجينات الكالسيوم مع الجيلاتين مشيرة الي الخلط المتجانس بينهما. وأثبتت خرائط EDX لجميع العينات تواجد اشعاع Kα لعناصر الكربون والاكسجين والكالسيوم والمفترض تواجدها في بوليمرات ألجينات الكالسيوم وتظهر النتائج ايضا في العينات المحتوية علي الهيدروكسي أباتيت ان الأطوار غير العضوية في قشور البيض وعظام الابقار تتكون اساسا من عناصر الكالسيوم والفوسفور بالاضافة لكميات قليلة من عناصر الكربون والاكسجين والصوديوم والماغنسيوم بالاضافة الي السيليكون. مع الحصول على نسبة مولارية لعنصر الكالسيوم مع الفوسفور (P/Ca) في الهيدروكسي أباتيت المحضرة معمليا مساوية 1.67 ولكن تزيد هذه النسبة بسبب تواجد كميات زائدة من كالسيوم الروابط البينية في الجينات الكالسيوم.
بالنسبة لعينات CAB و CAMK و CAD فيوجد اشعاع لعنصر السيليكون مُثبتاً مساهمة مركبات السيليكون في هذه المخاليط بنسب كبيرة والتي تساعد في التخلص من كاتيونات العناصر الثقيلة بدرجة كبيرة.
التعرف علي المجموعات الوظيفية من قياسات الاشعة تحت الحمراء بعد عملية الامتصاص (iv)
اشارت نتائج FTIR انه لايوجد تغير كبير في مواضع حزم الامتصاص بعد امتصاص ايونات المعادن وكذلك لم يتم الكشف عن اي ذٌرو امتصاص جديدة. ومع ذلك , هناك قليل جدا من الإزاحة لذروات الامتصاص , ويمكن ان يٌعزي ذلك الي اشتراك واحلال المعادن في الشبكة البلورية لألجينات الكالسيوم ومخاليطة المختلفة.
(v) قياسات حيود الاشعه السينية لعيينات للمركبات المحضرة بعد عملية الامتصاص
أجريت تحاليل الاشعة الشعة السينية XRD لعينات CA وCAHA(I) باعتبارهما الاعلي في عملية الامتصاص. تظهر النتائج انه لايوجد اطوار جديدة بعد امتصاص ايوني الرصاص الثنائي والحديد الثلاثي مما يدعم افتراض عدم حدوث الامتزاز نتيجة ميكانيكية التفكك والترسيب وانه ربما تحدث نتيجة عملية الامتزاز مثل التبادل الايوني وتكوين المتراكبات , وتظهر النتائج ايضاً حدوث تغيرات طفيفة في الشدة النسبية والحجوم البلورية وكذلك قيم التباعد d- spacing ويمكن ان يعزي ذلك الي التبادل الايوني بين ايونات المعادن و (1) ايونات الكالسيوم الموجودة في ألجينات الكالسيوم في CA و CAHA(I) كما في التفاعل التالي :
Ca(ALG)2 + Pb2+ Pb(ALG)2 + Ca2+
أو (2) ايونات الكالسيوم الخاصة بالهيدروكسي اباتيت في عينة ال CAHA(I) مكونا طور جديد من الهيدروكسي بيرومورفيت والتي لم تظهر في تحاليل XRD ربما لانها تكون بنسبة ضئيلة جدا تحت المدي الحثي XRD.
Ca10 (PO4)6 (OH)2 + x Pb2+ x Ca2+ + Ca10-xPbx (PO4)6(OH)2
(vi) التعرف على الشكل المورفولوجى والتحليل النوعي والكمى للعينات المحضرة باستخدام قياسات الميكروسكوب الإلكترونى الماسح للضوء وطاقة الأشعة السينية المشتتة بعد امتصاص ايونات الفلز.
كشفت النتائج عن بعض التغيرات في الشكل والبنية المجهرية ل CA و CAHA(I) عند التفاعل مع ايونات معادن الحديد والرصاص علاوة علي ذلك اشارت نتائج EDX الي وجود ايونات Pb2+ وFe3+ , كما تظهر النتائج أن ايونات الحديد الثلاثي هي اكثر كثافة من ايونات الرصاص الثنائي وهذا يوكد علي ان ايونات Fe3+ اكثر ازالة بواسطة CA و CAHA(I) مقارنة بايونات Pb2+. وتظهر النتائج ايضا انخفاض في نسبة ايونات الكالسيوم بعد عملية الامتصاص عنها قبل عملية الامتصاص, وهذا يمكن إيعازه الي احتمالية حدوث عملية الامتصاص نتيجة التبادل الايوني والذي يتوافق مع نتائج قياسات FTIR و تحاليل XRD. أيضا ربما تحدث عملية الامتصاص نتيجة ارتباط مخلبي لأيونات الفلزات مع مجموعتين كربوكسيليتين من الألجينات ومجموعة أواثنين من مجموعات الهيدروكسيل للألجينات , وفي هذه الحالة ربما يُكًون أيون الفلز متراكبات مع حلقتين متجاورتين لسلسلة بوليمرية واحدة من الألجينات او حلقتين لسلسلتين متوازيتين.
(vii) دراسة معدل امتصاص ايونات الفلزات تحت تاثير زمن التلامس
تمت دراسة تأثير الزمن علي قدرة ألجينات الكالسيوم ومخاليطة المختلفة في التخلص من ايونات الرصاص عند إس هيدروجيني للمحلول المحضر وهو pH=5.7 و عند 4pH = وايونات الحديد عند الاس الهيدروجيني للمحلول المحضر وهو 2.6 , وتشير الدراسة الي انه تم تحقيق نقاط الاتزان لجميع المخاليط في خلال من 10 الي 60 دقيقة من بداية الامتزاز ومراحل امتزاز (5-30) دقيقة , بنسب ازالة تتراوح ما بين66% إلي 83% للرصاص عند pH=5.7 و من 68% إلي 87% عند pH=4 وتتراوح مابين86% إلي 99% للحديد عند pH=2.6 .
وتكون عملية الامتزاز سريعة في المراحل الاولي ومتساوية تقريبا عند ازمنة التلامس العالية ويٌعزي هذا السلوك الي توافر عدد كبير من المواقع النشطة خلال المراحل الاولي من عملية الامتزاز وبالزيادة الكبيرة في زمن التلامس لم يُلاحظ حدوث أي عملية امتزاز وذلك بسبب الانخفاض في معدل الانتشار حيت ان جميع المواقع قد تم تغطيتها بايونات الفلز.
(viii) دراسة معدل امتصاص ايونات الفلزات تحت تاثير درجة الحموضة
تم دراسة تغيير درجة الإس الهيدروجيني علي عملية الامتزاز بالنسبة لايونات الرصاص فقط حيت ان ايونات الحديد Fe3+ تكون ثابته فقط عند اس هيدروجيني اقل من 3. تمت الدراسة عند قيمتين للإس الهيدروجيني وهما 5.7 (pH للمحلول المحضر) وعند pH=4, تركيز 100 ملجرام/لتر من ايونات الرصاص وكتلة من المادة المازّة ( (adsorbent 20 ملجرام/ 10 ملليتر من محلول ايون الفلز عند درجة حرارة 25 درجة مئوية. اظهرت الدراسة ان كفاءة عملية الإمتزاز عند 4 = pH تكون أعلي منها عند 5.7 = pH لجميع المخاليط.
(ix) دراسة معدل امتصاص ايونات الفلزات تحت تاثير كتلة الممتز
تم دراسة تاثير كتلة المادة المازّة بالنسبة لمعدل الامتصاص وذلك بالنسبة لإمتزاز الرصاص علي مركب ألجينات الكالسيوم (كنموذج قياسي لبقية المخاليط). تمت الدراسة باستخدام كتل مختلفة 10 و 15 و 20 ملجرام من المادة المازّة و10مليلترمن محلول الرصاص بتركيز 100 ملجرام/لتر و5.7 = pH ودرجة حرارة 25 درجة مئوية وازمنة تلامس من 5 الي 30 دقيقة. وتظهر النتائج ان معدل الامتصاص يزداد بزيادة كتلة المادة المازّة وهذ يٌعزي الي زيادة مساحة السطح ومن ثم زيادة المواقع النشطة المتاحة لعملية الامتزاز وزيادة كفاءة الازالة.
(x) دراسة كيناتيكية لعملية الإمتزاز
تمت دراسة حركية عملية الأمتزاز لمعدل امتصاص أيونات الفلزات علي سطح ألجينات الكالسيوم ومركباته المختلفة وتمت الدراسة عند ظروف تجريبية ( 25 درجة مئوية , كتلةمن المادة المازّة تساوي 20 مليجرام لكل 10 مليلتر من محاليل ايونات Pb2+ وFe3+ بتركيز ابتدائي 100 ملجرام/لتر) , وتم اختبار نموذجين حركيتين شائعتين تحت الظروف التجريبية وهما: معادلة الرتبة الأولى الكاذبة لـ Lagergren ومعائلة الرتبة الثانية الكاذبة لتحليل معدل امتزاز أبونات الفلزات علىCA و CACs.
1- نموذج الرتبة الأولى الكاذب
بتطبيق معادلة لاجرجرين لتفاعل الرتبة الاولى الزائفة التالية :
ln (qe –qt ) = ln qe – k1 t
حيث ان K1هو ثابت معدل لاجارجرين للإمتزاز (دقيقة-1) وqe وqt هى كميات العناصر الممتزة (ملجرام/جرام) عند الإتزان وعند الزمنt . ومن العلاقة البيانية بين Log(qe – qt) والزمن t, اتضح أن قيم معامل الارتباط R2تكون بعيدة وغير مقتربه من الوحدة كما ان قيم qe الناتجة من معادلة الرتبة الاولي الكاذبة تختلف تماما عن qe التجريبية , وهذا يشير الي أن معادلة Lagergren من الدرجة الأولى غير مناسبة لوصف امتزاز أيونات الفلزات المستهدفة بواسطة CA و CACs المستخدمة.
2- نموذج الرتبة الثانية الكاذب
بتطبيق معادلة تفاعل الرتبة الثانية الكاذبة التالية
t/qt = 1/k2qe2 + t/qe
على البيانات المعملية لإمتزاز ايونات الرصاص والحديد بواسطة ألجينات الكالسيوم ومركباته , حيث أن 2k هو ثابت معدل تفاعل الرتبة الثانية الزائفة (جم / ملجم. دقيقة) وqe وqt الكمية الممتزة في وحدة الكتلة عند الإتزان وعند الزمن t ، ومن العلاقة البيائية بين t/qt و الزمن t يتضح ان qe الناتجة من العلاقة الخطية لمعادلة الرتبة الثانية الكاذبة متوافقة تماما مع qe التجريبية , ووجد ايضا أن قيم معامل الارتباط تقترب جدا من الوحدة , وهذا يشير الي أن معادلة Lagergren من الدرجة الثانية مناسبة لوصف امتزاز أيونات الفلزات المستهدفة بواسطة CA و CACs المستخدمة
3 - تحديد الخطوة المتحكمة في التفاعل
من المعروف انه يوجد ثلاثة خطوات في أي عملية إمتزاز من محلول لمادة ممتزة وهي كالتالي (1) انتشار الطبقات الحدودية او الانتشار الفيلمي (2) انتشار بين الجزيئات او ثقبي (3) امتزاز علي المواقع او الاماكن الداخلية وهذه الخطوة سريعة جدا وربما لاتلاحظ , لذا فالخطوتين الاوليتين هما الخطوتين المتحكمتين في معدل الامتزاز. من اكثر الطرق المستخدمة في معرفة ميكانيكية عملية الامتزاز هي باستخدام الرسم البياني للانتشار بين الجزيئات للعالِمين ويبر وموريس بالعلاقة التالية:
qt = Kid t0.5 + C
حيث ان C ثابت و Kid هي ثابت الانتشار بين الجزيئات (ملجرام/جرام.دقييقة0.5 ( وqt هى كمية العناصر الممتزة (ملجرام/جرام) عند الإتزان. من العلاقة البيانية بين qt و t0.5 يتضح ان خطوة الانتشار بين الجزيئات ليست الخطوة المتحكمة في عملية الامتزاز لان العلاقة الخطية لا تمر بنقطة الاصل , كما توضح النتائج ان الامتصاص السريع للايونات في المرحلة الاولي يكون نتيجة الانتشار الفيلمي او تاثير الطبقات الحدودية وبعد اكتمال الأسطح الخاجية تتجه الأيونات للأنتشار بين الجزيئات اوخلال الثقوب.
(xi) دراسة الامتزاز عند درجة حرارة ثابته
من المفترض إن يكون إمتزاز ايونات العناصر + Pb2وFe3+ من الماء بواسطة CA و CAHA(I)
له سلوك يتلائم مع نموذج الإمتزاز عند ثبات الحرارة حيث أن المادة الممتزة تحافظ على الإتزان الديناميكي بين الإمتزاز وعدم الإمتزاز عند درجة حرارة ثابتة ويمكن تمثيل هذا النموذج بإستخدام معادلة لانجمير أو فريندلش. معادلة لانجمير لحساب أقصى قيمة لإمتزاز العناصر والتي تمثل بخط مستقيم تعطى من العلاقة
Ce/qe =Ce/qm + 1/ KL qm
حيثCe هو تركيز الإتزان للعنصر الثقيل المتبقي في المحلول(ملجرام/ لتر) عندما تمتز منه كمية تساوي qe. و qe هى الكمية الممتزة عند الإتزان ( ملجرام/ جم) وqm هى سعة الإمتزاز القصوى التي ترجع الى حدوث تغطية كاملة للسطح من طبقة واحدة (ملجرام / جم) و kLهوثابت لانجمير الذي يتناسب عكسيا مع طاقة الإمتزاز (ملجرام) ويمكن حساب qe من المعادلة
qe =((Co-Ce)V)/m
حيثCo هو تركيز الأيون الإبتدائي (ملجرام/ لتر) و Ct التركيز النهائي للعنصر (ملجرام/ لتر بعد مرور فترة من الزمن t , و V هو حجم المحلول الابتدائي (لتر) و m هي كمية الجينات الكالسيوم او مركبه مع الهيدروكسي اباتيت المضافة , كذلك من العلاقة التالية :
RL=1\1+ KLCe
نستطيع معرفة ما اذا كانت عملية الامتزاز مفضلة او غير مفضلة او غير عكسية او خطية حيث ان RL هي معيار الاتزان بلا ابعاد فاذا كانت قيمة RL تساوي صفر فان عملية الامتزاز تكون غير عكسية , واذا كانت تساوي 1 فان الامتزاز يكون خطي , واذا كانت بين صفر و واحد تكون مفضلة واذا كانت اكبر من الواحد تكون غير مفضلة.
اما معادلة فريندلش والتي تفترض وجود سطح غير متجانس مع توزيع غير منتظم لحرارة الامتزاز علي السطح وان عملية الامتزاز تتم علي طبقات عديدة و تمثل تلك المعادلة بخط مستقيم و تعطي من العلاقة
log qe =log KF + (1/n) log Ce
حيث KF و n هما ثوابت فريندلش ومن خلال قيمة (1/n) نستطيع معرفة ما اذا كان الامتزاز غير انعكاسي اذا كانت تساوي صفر اما اذا كانت بين 0 و 1 فان عملية الامتزاز تكون مفضلة اما اذا كانت قيمة (1/n) اكبر من 1 تكون عملية الامتزاز غير مفضلة.
من تطبيق تلك المعادلات علي النتائج التي تم الحصول عليها من عملية الامتزاز يتضح ان امتزاز ايونات الرصاص علي الجينات الكالسيوم يمكن تمثيلها بنموزج لانجمير بمعامل ارتباط جيد مايعني تغطية ايونات الرصاص لطبقه واحدة من CA , كذلك قيمة RL=0 تشير الي ان عملية الامتزاز تكون غير انعكاسية.
اما في حالة إمتزاز ايونات الرصاص والحديد علي سطح CAHA(1) وكذلك امتزاز ايونات الحديد علي سطح CA فيتضح أنها تُمثل جيدا بنموذج فريندلش بمعامل ارتباط جيد وكذلك قيمة (1/n) تشير الي ان عملية الامتزاز تكون مفضلة.
من خلال كل ماسبق من نتائج التحاليل المختلفة ومن دراسة كيناتيكية عملية الامتزاز يتضح أن:
إزالة عناصر الحديد والرصاص ربما تتم عن طريق تكوين متراكبات علي سطح ألجينات الكالسيوم ومركباته المختلفة او تتم عن طريق التبادل الايوني بين ايونات تلك العناصر وايونات الكالسيوم في عينات ألجينات الكالسيوم ومركباته المختلفة او ايونات الكالسيوم المرتبطة بالهيدروكسي اباتيت في عينات ألجينات الكالسيوم التي تحتوي علي مادة الهيدروكسي اباتيت.
مركب CAG هوالأعلي كفاءة في أزالة عنصر الرصاصPb2+ يليه مركبات CA و CAHA(I) ومركب CAHA(I) هو الأعلي كفاءة في إزالة عنصر الحديد Fe3+ يليه CAHA(II) ثم CA.
}
رسالة مقدمة من الطالب
محمد أحمد عبده عبدالله الامير
بكالريوس علوم (كيمياء – 2009)
كجزء من متطلبات الحصول علي
درجة الماجستير في العلوم
(كيمياء)
(كيمياء غير عضوية وتحليلية)
إلى
قسم الكيمياء
كلية العلوم
جامعة قناة السويس
الإسماعيلية
(2019)
دراسات علي إزالة العناصر الثقيله من المياه الملوثه باستخدام بعض المواد العضوية وغير العضويه المركبه
لجنة الأشراف التوقيع
1- أستاذ دكتور / صبري عبد الحميد القرشي ..........................
أستاذ الكيمياء غير العضوية
قسم الكيمياء
كلية العلوم
جامعة قناة السويس
2- دكتور / أيمن عبد المؤمن محمد مصطفي ..........................
مدرس الكيمياء الفيزيائية
قسم الكيمياء
كلية العلوم
جامعة قناة السويس
3- دكتور/ عباس ممدوح عباس ..........................
مدرس الكيمياء غير العضوية والتحليلية
قسم الكيمياء
كلية العلوم
جامعة قناة السويس
وافق مجلس الكلية بتاريخ / /
كما وافق السيد الأستاذ الدكتور / نائب رئيس الجامعة بتاريخ / /
علي منح درجة الماجستيرفي العلوم للطالب
محمد أحمد عبده عبدالله الامير
عنوان ارسالة:
دراسات علي إزالة العناصر الثقيله من المياه الملوثه باستخدام بعض المواد العضوية وغير العضويه المركبه
لجنة الحكم والمناقشة التوقيع
1- أستاذ دكتور / صبري عبد الحميد القرشي ..........................
أستاذ الكيمياء غير العضوية
قسم الكيمياء
كلية العلوم
جامعة قناة السويس
2- أستاذ دكتور / عصام عبدالعزيز ابراهيم كيشار ...........................
أستاذ الكيمياء غير العضوية
قسم الكيمياء
بنات عين شمس
3- أستاذ دكتور مساعد / خلود محمد ابو النور ...........................
أستاذ مساعد الكيمياء التحليلية
قسم الكيمياء
كلية العلوم
جامعة قناة السويس
وكيل الكلية لشئون الدراسات العليا عميد الكلية
أ.د/ علاء الدين عبدالعزيز سلام أ.د/ محمد سعد زغلول
Studies on Heavy Metals Removal from Polluted Water Using Some Organic and Inorganic Composite Materials
A Thesis Submitted by
Mohamed Ahmed Abdo Abd Allah Elamir
B.Sc in Chemistry
In Partial Fulfillment of the Requirements For
The Degree of master of Science (M.Sc.)
in
Chemistry (Inorganic and Analytical Chemistry)
to
Chemistry Departement
Faculty of Sience
Suez Canal University
Ismailia
(2019)
1. Introduction and Literature review
Heavy metals are metallic elements which have a comparatively high density compared to water. With the assumption that overweight and toxicity are interrelated, heavy metals also include metalloids, such as arsenic, that are able to motivate toxicity at low exposure level. Nowadays, the environmental pollution by these metals causes extreme ecological and global public health worry. Also, human exposure has risen intensely owing to an exponential increase of their use in various industrial, domestic, agricultural, and technological applications. Generally heavy metals sources involve industrial, geogenic, agricultural, pharmaceutical, domestic waste, and atmospheric sources. Environmental pollution is very notable in point source areas like mining, foundries and metal extraction, and other metal-based industrial operations [1-3].
Even though heavy metals are naturally occurring elements that are present throughout the earth’s crust, most environmental pollution and human exposure result from anthropogenic activities such as mining and smelting operations, industrial production and use, and domestic and agricultural use of metals and metal-containing compounds [3-5]. Environmental contamination can also occur through corrosion of metal, atmospheric deposition, soil erosion of metal ions and leaching of heavy metals, sediment resuspension, and metal evaporation from water resources to soil and groundwater. Natural phenomena such as weathering and volcanic eruptions have also been reported to significantly contribute to heavy metal pollution [1-3]. Industrial sources of heavy metals include metal processing in refineries, coal burning in power plants, petroleum combustion process, stations of nuclear power and high-tension lines, plastics, textiles, microelectronics, wood preservation, and paper-processing plants [6, 7]. It has been stated that metals such as iron, magnesium, manganese, chromium, molybdenum, cobalt, nickel, copper, selenium and zinc are necessary nutrients which are needed for various physiological and biochemical functions [1]. Insufficient supply of these micronutrients leads to a variety of deficiency diseases or syndromes.
Also, heavy metals are considered as trace elements owing to their presence in very small concentrations (ppb range to less than 10 ppm) in various environmental regions. Their bioavailability is affected by physical factors such as temperature, phase combination, adsorption, and sequestration. It is also influenced by chemical factors which affect speciation at thermodynamic equilibrium, kinetics of complexation, lipid solubility, and octanol/water partition coefficients. Biological factors, such as characteristics of species, trophic interactions, and biochemical/physiological adaptation, also play an important role [1, 8].
The essential heavy metals perform biochemical and physiological functions in plants and animals. They are important constituents of several key enzymes and play important roles in various redox reactions [1]. Copper, for example, serves as an essential cofactor for sundry oxidative stress-related enzymes including catalase, superoxide dismutase, peroxidase, cytochrome c oxidases, dopamine β-monooxygenase, monoamine oxidase, and ferroxidases [9-11]. Therefore, it is an essential nutrient that is incorporated into a number of metalloenzymes involved in hemoglobin formation, carbohydrate metabolism, catecholamine biosynthesis, and collagen cross-linking, elastin, and hair keratin. The ability of copper to cycle between an oxidized state, Cu(II), and reduced state, Cu(I), is used by cuproenzymes involved in reduction-oxidation reactions [9-11]. However, it is this characteristic of copper that also makes it potentially toxic as a result of the transitions between the two oxidation states Cu(II) and Cu(I) can lead to superoxide and hydroxyl radicals generation [9-12].
Additionally, too much exposure to copper causing cellular damage leading to Wilson disease in humans [11, 12]. As in the case of copper, several other essential elements are required for biologic functioning; however, an excess amount of this metals causes cellular and tissue damage leading to a variety of harmful effects and human diseases. For some elements including copper and chromium, a very small range of concentrations between useful and toxic effects is present . Other metals such as cadmium, antinomy, arsenic, aluminum, barium, beryllium, bismuth, gallium, germanium, gold, indium, lead, lithium, nickel, mercury, platinum, silver, strontium, tellurium, thallium, tin , titanium, vanadium, and uranium have no definite biological functions and are regarded as nonessential metals [1].
In biological systems, heavy metals have been reported to affect cellular organelles and components like lysosome, cell membrane, endoplasmic reticulum, nuclei, mitochondrial, and some enzymes involved in detoxification, damage repair, and metabolism [13]. Metal ions have been found to interact with cell components such as nuclear proteins and DNA, causing DNA damage and conformational changes that may lead to cell-cycle modulation, carcinogenesis, or apoptosis [13, 14]. Table (1) shows the main sources, health effects and permissible limits of various toxic heavy Metals.
Table (1): Sources, health effects and permissible limits of various toxic heavy Metals according to World Health Organization(WHO) [15, 16].
Metal
Source
Potential health effect potable water limits (ppm), WHO
Copper
Zinc
Mercury
Nickel
Cadmium
Arsenic
lead
Metal finishing industry
Electroplating
Metalliferous mining
Metal finishing industry
Electroplating, Fertilizers
Metalliferous mining
Agricultural material
Manures sewage sludge
Electronics
Waste disposal
landfill leachate
Metalliferous mining
Metal finishing industry
Electrodeposition
Manures sewage sludge
Alloys and steels
Metallurgical industries
Metalliferous mining
Agricultural materials
Fertilizers, waste disposal
Landfill leachate, electronics
Electronics
Metallurgical industries
Manures sewage
Specialist alloys and steels.
waste disposal
landfill leachate
Electronics, metallurgical
Industries
Nervous system irritation followed by depression
Liver damage, Wilson disease, insomnia
Phytotoxic, depression
Anemia, lethargy
Lack of muscular coordination
Abdominal pain
Increased thirst
Poisonous
Rheumatoid arthritis
Disturbs the cholesterol
diseases of the kidneys, circulatory system and nervous system
High conc. can cause DNA damage
Eczema of hands
Human carcinogen
High phytotoxicity
Damaging fauna
Kidney damage, renal disorder
Human carcinogen
Bronchitis, emphysema
Anemia
Acute effects in children
Skin manifestations,
visceral cancers, vascular disease
Immunotoxic
Modulation of co-receptor expression
Damage the fetal brain
diseases of the kidneys
circulatory system and
nervous system
2
3
0.001
0.02
0.03
0.01
0.01
1.1. `Mechanism of heavy metals in human being.
In past decades, several studies have been carried out to investigate the mechanism of toxicity with heavy metals [15] . Toxicity and carcinogenicity of heavy metals involve many mechanistic aspects, some of which are not clearly elucidated or understood [1].
Many studies have reported that the production of reactive oxygen species (ROS) and oxidative stress play an important and a key role in the toxicity and carcinogenicity of metals for example arsenic [17-19], cadmium [20] , chromium [21, 22] , lead [23, 24], and mercury [25, 26] . For the reason of their high degree of toxicity, these previous metals are among the priority metals that are of great public health significance. They are all systemic toxicants that are known to prompt multiple organ damage, even at lower levels of exposure. As per the US Environmental Protection Agency (US EPA) and the International Agency for Research on Cancer (IARC), these elements are also categorized as either “known” or “probable” human carcinogens based on epidemiologicalds and experimental studies showing a link between exposure and cancer incidence in humans and animals.
Oxidative stress is one of the major mechanisms behind metal toxicity [27]. The formation of large amounts of reactive oxygen species, such as superoxide anion (O2.-), hydrogen peroxide (H2O2), hydroxyl radical (HO.) and singlet oxygen (1O2), has been reported to promote the induction of oxidative stress. [15, 28] In fact, various studies connect heavy metals with oxidative DNA damage since these metals may reduce the level of the main antioxidant compounds in several animal tissues by inactivating enzymes and other antioxidant molecules [29]. In humans, oxidative stress is also responsible for various diseases, including cancer, Parkinson’s disease, Alzheimer’s disease, atherosclerosis, heart failure and myocardial infarction [30, 31].
Although Heavy metal-induced toxicity and carcinogenicity involve many mechanistic aspects, some of which are not clearly elucidated or understood. Each metal is known to have unique features and physicochemical properties that confer to its specific toxicological mechanisms of action [1].
1.1.1 BIOCHEMISTRY OF TOXICITY
The heavy metals poisoning effects are because of their interference with the normal body biochemistry in the normal metabolic processes. When this metals ingested, in the acid medium of the stomach, they are converted to their stable oxidation states (Zn2+, Pb2+, Cd2+, As2+, As3+, Hg2+ and Ag+) and combine with the body’s biomolecules such as proteins and enzymes to form stable and strong chemical bonds. The equations shown in Figure (1) display their reactions during bond formation with the sulphydryl groups (-SH) of cysteine and sulphur atoms of methionine (-SCH3) [32, 33].
Figure (1): interaction between metal with proteins and enzymes.
(A) = Intramolecular bonding; (B) = Intermolecular bonding; P = Protein; E = Enzyme; M = Metal
The metal groups or the hydrogen atoms in the above case are substituted by the poisoning metal and the enzyme is thus inhibited from functioning, whereas the protein–metal compound works as a substrate and can reacts with a metabolic enzyme. In the following scheme, equation C indicates the reaction of enzymes (E) with substrates (S) in either the lock-and-key pattern or the induced-fit pattern. In both cases, a substrate fits into an enzyme in a highly specific fashion, as aresult of enzyme chirality’s, to form an enzyme–substrate complex (E-S*) as follows [33] .
(E = Enzyme; S = Substrate; P = Product; * = Activated Complex)
While at the E-S, E–S* and E-P states, an enzyme cannot accommodate any other substrate till it is freed. Occasionally, the enzymes for an entire sequence coexist together in one multi-enzyme complex consisting of three or four enzymes. The product from one enzyme reacts with a second enzyme in a chain process, with the last enzyme yielding the final product as follows:
The final product (F) goes back to react with the first enzyme thereby inhibiting further reaction since it is not the starting material for the process. Hence, the enzyme E1 becomes incapable of accommodating any other substrate until F leaves and F can only leave if the body utilizes it. If the body cannot utilize the product formed from the heavy metal – protein substrate, there will be a permanent blockage of the enzyme E1, which then cannot initiate any other bio-reaction of its function. Therefore, the metal remains embedded in the tissue, and will result in bio-dysfunctions of various gravities. Furthermore, a metal ion in the body’s metallo-enzyme can be conveniently replaced by another metal ion of similar size. Thus Cd2+ can replace Zn2+ in some dehydrogenating enzymes, leading to cadmium toxicity. In the process of inhibition, the structure of a protein molecule can be mutilated to a bio-inactive form, and in the case of an enzyme can be completely destroyed. For example, toxic As3+ occurs in herbicide, fungicides and insecticides, and can attack –SH groups in enzymes to inhibit their bioactivities as shown below in Figure (2) [32, 33].
Figure (2): interactions between arsenic and enzyme.
The most toxic forms of these metals in their ionic species are the most stable oxidation states. For example, Cd2+, Pb2+, Hg2+, Ag+ and As3+. In their most stable oxidation states, they form very stable biotoxic compounds with the body’s bio-molecules, which become difficult to be dissociated, due to their bio-stabilities, during extraction from the body by medical detoxification therapy.
1.2. Heavy metals treatment techniques:
Heavy metal uptake from inorganic effluent and industrial wastewaters can be carried out by traditional treatment processes, such as, complexation, adsorption, coagulation, ion exchange, solvent extraction, chemical precipitation, electroplating, cementation, flotation and membrane separation. All of these processes may be physical, chemical or biological as shown in Figure (3), Some of these are illustrated in Figure (4) [34] . Different methods, such as chemical precipitations, conventional adsorption [35-37] , ion exchange [38], membrane separation techniques [39] and electro-remediation techniques are used usually for industrial wastewater treatment. Precipitation is most economical and hence widely used, but many industries still use chemical procedures for treatment of effluents due to economic considerations. The efficiency of the precipitation process is extremely decrease owing to the presence of complexing agents in wastewater , and this lead to incomplete processing and production of toxic sludge. Thus several new approaches have been studied to develop low cost effective and more efficient heavy metal adsorption techniques [40].
Biosorption is considered as a user-friendly with specific affinity, low cost and simple design so it is an effective separation and purification method for heavy metals disposal from industrial wastewater and it has been widely used for this purpose [41, 42].
Figure (3): Conventional technologies for heavy metal removal.
Sorption with sorbents made of agricultural or industrial by-products are used widely for heavy metals uptake from aqueous mediums because of their abundant availability, promising physical, low cost, and, surface and chemical characteristics [43]. Those materials and methods were widely discussed meeting their advantages.
Figure (4): Some conventional methods for metal removal.
1.2.1. Physico-chemical methods
Following methods have been used by various researchers for heavy metals uptake. Physical separation process are primarily applicable to particulate forms of metals, metal-bearing particle or discrete particles . Physical separation consists of flotation, mechanical screening, gravity concentration, hydrodynamic classification, magnetic separation, attrition scrubbing, and electrostatic separation, physical separation efficiency depends on several soil characteristics such as particulate shape, particle size distribution, moisture content, humic content, clay content, density between soil matrix ,metal contaminants and heterogeneity of soil matrix, and hydrophobic properties, magnetic properties of particle surface [40, 44].
The conventional chemical processes for heavy metals disposal from waste water contain many processes such as flotation, ion exchange, adsorption, electrochemical deposition and chemical precipitation. Factors which may limit the effectiveness and applicability of the chemical process are high content of clay/silt, calcite, humic, Ca and Fe, anions, heavy metals, or high buffering capacity [45].
A. Chemical Precipitation:
Chemical precipitation is one of the most widely used for removal of heavy metal from inorganic effluent in industry because of its simple operation[46]. These conventional chemical precipitation processes yield insoluble precipitates of heavy metals as hydroxide, carbonate, phosphate and sulfide. The mechanism of this process is depend on to produce insoluble metal precipitation by reacting dissolved metals in the solution and precipitant. In the precipitation method very fine particles are generated and chemical precipitants, flocculants and coagulation processes are used to increase their particle size in order to remove them as sludge [45, 46]. As soon as the metals precipitate and form solids, they can easily be removed, and metal with low concentrations, can be discharged. Removal percentage of metal ions in the solution may be improved to optimum by changing major parameters such as pH, initial concentration, ions charge, temperature,.…etc. Hydroxide treatment is the most commonly used precipitation technique owing to its relative simplicity, low cost of precipitant (lime), and ease of automatic pH control. The solubilities of the various metal hydroxides are minimized for pH in the range (8-11).
B. Coagulation and Flocculation:
The coagulation-flocculation mechanism is based on zeta potential (ζ) measurement as the standard to define the electrostatic interaction between coagulant-flocculant agents and pollutants [47] .Coagulation process is reduced the net surface charge of the colloidal particles to stabilize by electrostatic repulsion process [40]. Flocculation process continually increases the particle size to discrete particles through additional collisions and interaction with inorganic polymers formed by the organic polymers added [48]. The minute discrete particles are flocculated into larger particles, they can be removed or separated by filtration, floatation or straining. Sludge production, transfer of toxic compounds into solid phase and application of chemicals are main disadvantages of this process.
C. Electrochemical Treatments:
Electrolysis: Electrolytic recovery is one technology used for removing metals from waste water streams. This process uses electricity to pass a current through an aqueous metal-bearing solution containing a cathode plate and an insoluble anode. Electricity can be generated by movements of electrons from one element to another. Electrochemical process to treat wastewater containing heavy metals is to precipitate the heavy metals in a weak acidic or neutralized catholyte as hydroxides. Electrochemical treatments of wastewater involve electro-deposition, electro-coagulation, electro-flotation and electro-oxidation [49].
Electro-destabilization of colloids is called coagulation and precipitation by hydroxide formation to acceptable levels. It is the most common heavy metal precipitation process forming coagulants by electrolytic oxidation and destabilizing pollutants to form folc [50]. The electro-coagulation process the coagulant is generated in situ by electrolytic oxidation of an appropriate anode material. In this process, charged ionic metal species are removed from wastewater by allowing it reacting with anion present in the effluent. This process is characterized by reduced production of sludge, ease of operation and no requirement for chemical use.
However, chemical precipitation requires a large amount of chemicals to reduce metals to permissible limit for discharge. Other drawbacks are huge sludge production, poor settling, slow metal precipitation, long-term environmental impacts of sludge disposal , and the aggregation of metal precipitates [51]. It converts the aqueous pollution problem to a solid waste disposal problem without recovering the metal.
D. Ion Exchange:
Ion exchange can attract soluble ions from the liquid phase to the solid phase. It considered is the most widely used technique in water treatment industry. As a cost-effective method, ion exchange process usually involves convenient operations and materials with low-cost, and it has been demonstrated to be very effective for elemination heavy metals from aqueous mediums, specific for treating water with low heavy metals concentration [52, 53]. In this technique cations or anions containing special ion exchanger is used to eliminate metal ions from the solution. Commonly used ion exchangers are synthetic organic ion exchange resins. It can be used only low concentrated metal solution and this method is extremely sensitive with the pH of the aqueous phase.
Ion exchange resins are water-insoluble solid substances which can absorb negatively or positively charged ions from an electrolyte solution and release other ions with the same charges into the solution in an equivalent amount. The positively charged ions in cationic resins such as sodium and hydrogen ions are replaced with positively charged ions, such as, copper, zinc and lead ions, in the solutions. In a similar way, the negative ions in the resins such as hydroxyl and chloride ions can be exchanged by the negatively charged ions such as nitrate, chromate, sulfate, cyanide, and dissolved organic carbon (DOC).
E. Membrane Filtration:
Membrane filtration has received a great attention for the remidation of inorganic effluent. It is able to remove organic compounds, suspended solid, and inorganic pollutants such as heavy metals. Depending on the particle size that can be retained, different types of membrane filtration such as nanofiltration, ultrafiltration, and reverse osmosis can be utilized for heavy metal uptake from wastewater.
Ultrafiltration (UF) utilizes permeable membrane to separate heavy metals, macromolecules and suspended solids from inorganic solution on the basis of the pore size ranging from 5 to 20 nm and molecular weight of the isolating compounds (1000– 100,000 Da) [54]. Based on the membrane characteristics, UF can achieve more than 90% of removal efficiency with a metal concentration (10 - 112 mg/L) at pH ranging between 5 and 9.5 and at pressure (2–5 bar) . UF offers some advantages such as smaller space requirement and a lower driving force owing to its high packing density.
Polymer-supported ultrafiltration (PSU) process adds water soluble polymeric ligands to bind metal ions and form macromolecular complexes by generating a free targeted metal ions effluent [55]. PSU technology has the dvantages of the low-energy requirements involved in ultrafiltration, the reaction kinetics is very fast and higher separation selectivity of selective bonding agents in aqueous solution.
Another similar technique, complexation–ultrafiltration, confirms to be a promising alternative to technologies depend on ion exchange and precipitation. Using water-soluble metal-binding polymers in combination with ultrafiltration (UF) is a hybrid approach in order to concentrate selectively and to recover heavy metals in the solution. In the complexation – UF process cationic forms of heavy metals are first complexed by a macro-ligand in order that increasing their molecular weight with a size larger than the selected membrane pores [56, 57]. The advantages of complexation–filtration process iclude high selectivity of separation because of the use of a selective binding and low-energy requirements involved in these processes. Water-soluble polymeric ligands have shown to be powerful substances in order to separate trace metals from aqueous solutions and industrial wastewater through membrane processes.
Reverse osmosis (RO) is a separation process using pressure in which solution is forced through a membrane that keeps the solute on one side and allows the passage of the pure solvent to the other side. The membrane here is semi-permeable, meaning it allows the passage of solvent but not for metals. The reverse osmosis membranes have a dense barrier layer in the polymer matrix where most separation takes place. Reverse osmosis can remove many types of ions and molecules from solutions, including bacteria, and is used in both industrial processes. A diffusive mechanism is involved in reverse osmosis process, so that separation efficiency is dependent on pressure, concentration of the solute, , and water flux rate [58].
F. Electrodialysis:
Electrodialysis (ED) is a membrane separation uses electric potential to passe ionized species in solution through an ion exchange membrane. The membranes are plastic materials thin sheets with either cationic or anionic characteristics. When a solution containing ionic species passes through the compartments of the cell, the anions migrate toward the anode while the cations migrate toward the cathode, crossing the anion exchange and cation-exchange membranes [59]. A disadvantage of this process include membranes replacement and the corrosion process [60]. Using membranes with higher capacity of ion exchange resulted in better cell performance. Effects of temperature, flow rate,and voltage at different concentrations by using two types of commercial membranes, using a laboratory ED cell, on the removal of lead were studied [61]. The princible of Electrodialysis process is illustrated in Figure (5). Results show that increasing temperature and voltage improved cell performance and separation percentage decreased as the flow rate increasing. This provides advantages for the treatment of highly concentrated wastewater laden with heavy metals to recovery undesirable impurities from water.
Figure (5): Electrodialysis principles [62] CM – cation exchange membrane, D-dialute chamber, e1 and e2-electrode chambers, AM-anion exchange membrane and K-concentrate chamber
G. Adsorption:
Biosorption is another technique that can used for elimination of heavy metals from wastewater. Sorption process is defined as transfer of ions from solution phase to the solid phase, really describes a group of processes, which includes adsorption and precipitation reactions. Adsorption has become one of the alternative remidation techniques for wastewater. Basically, adsorption is a mass transfer process and substances bound by chemical and or physical interactions to solid surface [63-65]. All adsorption mechanisms are dependent on solid-liquid equilibrium and on mass transfer rates. A dsorption could be divided into the following types, depending on the types of intermolecular attractive forces [66, 67].
Physical adsorption:
It is a process in which binding of adsorbate on the surface of adsorbent as a result of Van der Waals forces of attraction or hydrogen bonding. Physical adsorption can only be occurred in the low temperature environment and under appropriate pH conditions.
Chemical adsorption:
A strong interaction arise from chemical reaction between the adsorbate and the adsorbent molecules is involved. This interaction produces new types of electronic bonds (Covalent and Ionic).
Mechanism of adsorption:
generally, the main steps involved in adsorption of contaminates on solid adsorbent are:
1.Transfer of the metal ion from the bulk of solution to the outer surface of the adsorbent.
2. Internal mass transfer by pore diffusion from outer surface of adsorbent to the inner surface of porous structure.
3. Adsorption of adsorbate onto the active sites of the adsorbent pores
4. The overall adsorption rate is determined by either intra particle diffusion or film formation or both as the last step of adsorption are very fast as compared to the other two steps.
The parameters which have been established for optimizing the use of adsorbent in wastewater treatment include [34]:
1. Nature of adsorbent and adsorbate.
2. Metal concentration.
3. pH and temperature of the aqueous solution.
4. Kinetics of adsorption.
5. Adsorption isotherm.
6. The time of contact.
Various low-cost adsorbents, derived from natural material, agricultural waste, industrial by-product, or modified biopolymers are found to be more promising and encouraging in heavy metal removal owing to various considerations as follow [36, 63].
(I)They are economical, (II) its metal selectivity, (III) they are regenerative, (IV) toxic production of sludge not present (V) metal recovery and (VI) its high effectiveness.
Using activated carbon in water and wastewater remidation has been directed towards organics removal [40] . Research efforts on removal of inorganics by activated carbon, specifically metallic ions, have been markedly limited [40] . selective adsorption by red mud [68], coal [69], photocatalyst beads [70], nano-particles[71], fertilizer industrial waste [72], biomass [73], activated sludge biomass [74], algae [75, 76] etc. has generated increasing excitement.
Industrial by-products such as fly ash [77] iron slags, waste iron [78] , hydrous titanium oxide [79, 80] ,can be chemically modified to enhance its removal performance for metal elimination from wastewater.
It was reported that [81, 82] for the disposal of heavy metals from industrial waste effluent has been focused on the use of agricultural by-products as adsorbents through biosorption process. New resources such as rice husk, coconut shell, pecan shells, rice straw, maize cob or husk, jackfruit, hazelnut shell, rice husk,…etc can be used as an adsorbent after chemical modification or conversion by heating into activated carbon or biochar for heavy metal uptake. They found that the maximum metal removal occurred by those biomass due to containing of cellulose, lignin, carbohydrate and silica in their adsorbent [83] .
Biopolymers are posse a number of different functional groups, such as amines and hydroxyls, which increase the efficiency of metal ion uptake [84] . They are widely use in industrially as they are able to lower the concentrations of transition metal ion to sub-part per billion concentrations. New polysaccharide-based-materials are described as biopolymer adsorbents (derived from chitosan, starch and chitin) for the elimination of heavy metals from the wastewater. The sorption mechanisms of polysaccharide-based-materials are complicated and depend on pH [84]. Also hydrogels, which are cross linked hydrophilic polymers, are widely used to purify wastewater. The removal is mainly governed by the water diffusion into the hydrogel, carrying the heavy metals inside especially in the nonexistence of strongly binding sites. Maximum binding capacity increases with higher pH because of polymerization/cross linking reaction.
1.2.2. Biological Methods:
Biological removal of heavy metals in wastewater involves the use of biological methods for the elimination of pollutants from wastewater. In this processes microorganisms play an important role of settling solids in the solution. Activated sludge, stabilization ponds, trickling filters are widely used for wastewater purification. Activated sludge is considered the most common option uses microorganisms in the treatment process to break down organic matter with agitation and aeration, and then allows solids to settle out. Bacteria-containing “activated sludge” is frequently re-circulated back to the aeration basin to increase organic decomposition rate. In biological systems, most of the research on heavy metals removal has been oriented towards the suspended growth activated sludge process. Trickling filters which consist beds of coarse media (often plastic or stones) 3-10 ft. deep help to grow microorganisms. Wastewater is sprayed into the air (aeration), then allowed to trickle through the media and microorganisms break down organic matters in the wastewater. The drain of trickling filters at the bottom and the wastewater is collected and then undergoes sedimentation. Lagoons or stabilization ponds are cheap, slow and relatively inefficient, biological method that can be used for different types of wastewater. They depend on the interaction of sunlight, microorganisms, algae, and oxygen [40].
1.3. Evaluation of heavy metals removal processes:
Although all the techniques of heavy metal wastewater treatment can be employed to eliminate heavy metals, they have their latent advantages and limitations. Table (2) indicates heavy metals uptake from aqueous solutions has been traditionally carried out by chemical precipitation because it is a simple process and cheap capital cost. However, chemical precipitation is ordinarily adapted for treating wastewater containing heavy metal ions with high concentration and it is ineffective with low metal ion concentration. Chemical precipitation is considered as not economical and can produce large sludge amount to be treated with great drawbacks [85].
Ion exchange has been commonly applied for the removal of heavy metal from wastewater. However, the resins of ion-exchange must be regenerated by chemical reagents when they are exhausted and the regeneration can cause serious secondary contamination. And it is expensive, particularly when treating a large amount of wastewater containing low concentration heavy metal, so that they cannot be used at large scale.
Adsorption is a common method for the uptake of heavy metals from low concentration aqueous solutions containing heavy metal. The activated carbon high cost limits its use in adsorption. Many varieties of adsorbents with low-cost have been developed and tested for heavy metal ions uptake. However, the efficiency of adsorption depends on the adsorbents type. Biosorptio of heavy metals from aqueous mediums is considered as new method that has proven very promising for the removal of heavy metal ions from wastewater [85].
The technology of membrane filtration can separate heavy metal ions with high efficiency, but its drawbacks such as high cost, low permeate flux process complexity, and membrane fouling have limited their use in heavy metal ions uptake.
Coagulation-flocculation technique can be employed for heavy metal wastewater remidation, the advantages of this method are dewatering and good sludge settling of the produced sludge. large sludge volume generation and chemical consumption are the limitations of this technique.
Table (2): The main advantages and disadvantages of the various physico-chemical methods for treatment of heavy metal in wastewater.
Treatment method
Target of removal
Advantages
Disadvantages
References
Chemical precipitation
Coagulation–flocculation
Dissolved air flotation
Ion exchange
Ultrafiltration
Nanofiltration
Reverse osmosis
Adsorption with new adsorbents
Heavy metals, divalent metals
Heavy metals and suspended solids.
Heavy metals and suspended solids
Dissolved compounds,
cations/anions
High molecular weight compounds (1000–10000 Da)
Hardness ions such as Ca(II) and Mg(II) and sulphate salts
Organic and inorganic compounds
Heavy metals
Low capital cost, simple operation
Shorter time to settle out suspended solids, improved sludge settling.
Low cost, shorter hydraulic retention time
No sludge production, less time consuming
Smaller space requirement
Lower pressure than RO (7–30 bar)
High rejection rate, able to withstand high temperature
Low-cost, easy operating conditions, having wide pH range, high metalbinding capacities
Sludge generation, extra operational cost for sludge disposal.
Sludge production, extra operational cost for sludge disposal.
Subsequent treatments are required to improve the removal efficiency of heavy metal
Not all ion exchange resin is suitable for metal removal, high capital cost
High operational cost, prone to membrane fouling
Costly, prone to membrane fouling
High energy consumption due to high pressure required (20–100bar), susceptible to membrane fouling
Low selectivity, production of waste products
[86-89]
[87]
[87, 90]
[91, 92]
[91, 93]
[87]
[87, 91]
[81, 94, 95]
. Flotation presents several advantages over the more conventional methods, such as high removal efficiency, high selectivity of metal ions, low detention periods, low operating cost production of more concentrated sludge and high overflow rates [96]. Operation costs, high initial capital cost and high maintenance are the disadvantages of this process.
Electrochemical heavy metal wastewater treatment technologies are considered as rapid and well-controlled that require fewer chemicals, offer good reduction yields and generate less sludge. On the other hand, electrochemical methods involving high cost electricity supply and high initial capital investment, this restricts the technique development.
Biological technologies by using different low materials were found be very effective techniques with higher uptake percentage. Although biological techniques are low cost and friendly methods for the environment they require large areas and proper operation and maintenance [40].
Even though all above techniques can be employed for the remediation of heavy metal wastewater, it is important to remarkable that the selection of the most suitable remidation methods depends on the initial concentration of the metal, wastewater component, plant flexibility and accuracy, capital investment, operational cost and environmental impact, ...etc [60, 85].
1.4. Enviromental pollution with iron metal and its removal:
Iron is considered the second metal among the most abundant metals on the earth crust. In the periodic table of elements, iron occupies the 26th elemental position. Iron is existing in many forms in water as shown in Figure (6) [97]. Biologically it is a most crucial element for survival and growth of almost all living organisms. As it is the cofactor for many vital enzymes and proteins. It is one of the vital constituents of organisms like algae and of enzymes such as catalase and cytochromes, in addition to oxygen transporting proteins, such as myoglobin and hemoglobin. Because of iron inter-conversion between ferrous (Fe2+) and ferric (Fe3+) ions, it is regarded as an attractive transition metal for various biological redox processes owing. The iron source in surface water is anthropogenic and is associated with mining activities. Sulphuric acid production and the discharge of ferrous (Fe2+) occurs because of iron pyrites (FeS2) oxidation that are common in coal seams. [98-100]. The following equations represent the simplified oxidation reaction for ferrous and ferric iron [99]:
2FeS2 + 7O2 + 2H2O 2FeSO4 + 2H2SO4 (ferrous)
4FeSO4 + O2 + 10H2O 4Fe(OH)3 + 4H2SO4 (ferric)
Mediated reactions of iron support the respiration process of most of the aerobic organisms. If it is not shielded correctly, it can catalyze the reactions involving radicals formation which can destroy biomolecules, tissues, cells and the whole organism. Iron poisoning has always been a subject of interest chiefly to pediatricians. Children are highly susceptible to iron toxicity as they are exposed to a maximum of products containing iron [101].
Figure (6): Classification of different forms of Iron presen in water.
Iron toxicosis occurs in four stages [100]:
The first stage which takes place after 6 hrs of iron overdose is noticeable by effects of gastrointestinal such as vomiting, diarrhea and gastro intestinal bleeding.
The second stage progresses within (6 - 24hrs) of overdose and it is regarded as the inherent period, a period of apparent medical recovery.
The third stage happens between 12 to 96 hrs after certain clinical symptoms onset. This stage is characterized by shocks, tachycardia, lethargy, hypotension, metabolic acidosis hepatic necrosis, and sometimes death.
The fourth stage take place in betwwen 2 to 6 weeks of iron overdose. This stage is marked by the gastrointestinal ulcerations formation and strictures development.
Iron uptake excess is a serious problem in meat eating and developed countries and it increases the cancer risk. Workers who are highly susceptible to asbestos that contains almost 30% of iron are at high risk of asbestosis, which is the second most important reason for lung cancer. It is said that asbestos associated cancer is related to free radicals. Loose intracellular iron can also promote DNA destruction. Iron can initiate cancer mainly by the DNA oxidation process.
Iron salts such as iron sulfate, iron sulfate heptahydrate and iron sulfate monohydrate are of low acute toxicity when exposure is through dermal, oral and inhalation routes and hence they have been placed in toxicity category 3. Moreover, the Food and Drug Administration considered that iron salts are safe and their toxic effects are very much negligible.
Free radicals formation is the outcome of the toxicity of iron. During pathological and normal cell processing, byproducts such as hydrogen peroxide and superoxide are produced, which are regareded to be free radicals. These free radicals are actually neutralized by enzymes such as catalase superoxide dismutase, and glutathione peroxidase but the superoxide molecule has the capability to release iron from ferritin and that free iron reacts with more and more of hydrogen peroxide and superoxide forming free radicals with high toxicity such as hydroxyl radical. Hydroxyl radicals are dangerous as they can initiate lipid peroxidation, inactivate certain enzymes, cause DNA strand breaks and can depolymerize polysaccharides. This can occasionally lead to cell death.[100, 102, 103]
Tahir and Rauf [104] studied the Removal of Fe(II) from the galvanized pipe manufacturing industry wastewater by adsorption onto bentonite clay. The adsorption of Fe(II) from aqueous solutions over a concentration range from 80 to 200 mg/l, shaking time of 1–60 min, adsorbent dosage 0.02 – 2 g and pH of 3. The process of removal follows both the Langmuir and Freundlich isotherm models and also obey the first-order kinetics. The maximum removal (> 98%) was observed at pH of 3, 0.5 g of bentonite with initial concentration of 100 mg/l. The Fe(II) removal efficiency was also tested using wastewater from a galvanized pipe manufacturing industry. Higher than 90% of Fe(II) can be effectively removed from the wastewater by using 2.0 g of the bentonite. The effect of cations (i.e. manganese, cadmium, lead, chromium, zinc, copper, nickel and cobalt) on the removal of Fe(II) was studied in the concentration range of 10–500 mg/l. All the added cations reduced the Fe(II) adsorption at high concentrations except Zn. Column studies have also been investigated using a certain concentration of wastewater. More than 99% recovery has been attained by using 5 g of the bentonite with nitric acid solution (3 M).
Das et al. [105] in a study investigated that the traditional method of using ash for disposal of iron from groundwater can eliminate iron to desired level without increasing the pH behind the acceptable limit. The banana pseudostem ash is among the different plant ashes used for iron removal. It has been found to be most appropriate for iron elemination. The ash improves iron uptake. The designed iron elemination system is expected to be convenient for household use. The optimum values of the different parameters for iron removal are 200–300 mg L−1 ash, 1.0 L h−1 rate of filtration and time of residence (1h) for groundwater having 2.20 ppm iron concentration. For groundwater having higher [Fe], the amount of ash can be increased and can be decreased gradually throughout continuous use. The technique has the advantages of low manufacturing cost, almost nil recurring cost, viz., simplicity in use, no electricity requirement, and increasing the essential minerals such as K, Ca in the treated water.
Al-Anber et al. [106] in their study the batch removal of Fe3+ from aqueous model solution under different experimental conditions using Jordanian natural zeolite (JNZ) has been investigated. The contact time influences, initial concentration of metal, temperature and concentration of adsorbent dosage have been studied. The adsorption efficiencies are found to be residence time dependent, increasing the contact time in the range between 1 and 150 min. The sorption equilibrium has achieved between 60 and 150 min. The optimum adsorption has occur at 30°C of temperature. The equilibrium adsorption capacity of JNZ adsorbent used for Fe3+ were evaluated and extrapolated using Freundlich and Langmuir isotherm models and the experimental data are found to fit Langmuir isotherm more than Freundlich isotherm.
Ghosh and his team [107] reported the results of astudy on electrocoagulation (with electrodes from aluminum) for iron Fe(II) elemination from aqueous medium. The removal of Fe(II) was composed of two principal steps; (a) oxidation of Fe(II) to Fe(III) and (b) subsequent Fe(III) disposal by the freshly formed aluminum hydroxides complexes by adsorption/surface complexation followed by precipitation. Experiments were executed with various current densities ranging between 0.01 and 0.04 A/m2. Other parameters such as salt concentration, pH and conductivity were maintaned constant as per tap water quality. It was observed that as the current densities increase, the elemination of Fe(II) increased. Satisfactory iron removal of around 99.2% was attained at the end of 35 min of operation from 25 ppm initial Fe(II) concentration.
Vasudevan et al [108] reported the results of astudy on the uptake of Iron from drinking water by electrocoagulation using galvanized iron as the cathode and magnesium as the anode. Experiments were done as a function of current density, pH and temperature. By using both the Langmuir and the Freundlich isotherm models, the adsorption capacity was estimated. The results demonstrated that the maximum efficiency of removal of 98.4% was obtained at 0.06 A dm– 2 of current density and pH of 6.0. The adsorption of iron was better illustrated by fitting the Langmuir adsorption isotherm, which suggests adsorbed molecules monolayer coverage. The adsorption process followed a kinetics model of second-order. Temperature studies illustrate that adsorption was endothermic and spontaneous in nature.
Bulai and Cioanca [109] in astudy found that Purolite S930 is an effective sorbent for Iron (II) ions uptake from aqueous model solutions in different operating conditions. The Iron (II) elemination percent has a maximum at pH 5.0 and increases with contact time and resin dose increasing and decreases with solution initial concentration increasing.
Wang [110] investigated on crushed concrete and limestone removed Fe(II) from synthetic groundwater in laboratory columns. The results shown that achieving average Fe(II) uptake of greater than 216 (approximately 133 L treated) and 99% over 288 (approximately 172 L treated) pore volumes, for crushed concrete and limestone, respectively. Calcium siderite which formed in limestone columns as a form of precipitate; this formation had no significant effect on porosity of the system, but may have impeded Fe(II) remidation by limiting available surface area for adsorption. Results suggested that field-scale passive iron removal systems, using similar materials,merit exploration. Because differences are expected, pilot-scalefield tests are warranted.
Nandeshwar et al. [111] in astudy foucsed on iron uptake from real wastewater samples of Nag River, India using Green activated carbons from different waste materials such as orange peels, sawdust, C. procera leaves and coconut shells. All the selected waste materials were carbonized in muffle furnace and activated using various agents such as HCl, HNO3, and H2SO4. The results showed that all adsorbents have the potential capacity to separate iron, which further highly increases after its activation. The most promising green adsorbents were found to be orange peels and HCl was the best activating agent. The order of iron uptake from wastewater is: orange peels then coconut shells then sawdust then C. procera leaves. Similarly it was found that charcoal activated with HCl can separate around 77–90% iron followed by HNO3 (70–80%) and H2SO4 (58–75%).
Vries et al. [112] conducted astudy on iron separation from water by using rapid sand filtration. A model has been developed that takes into account the main properites of (submerged) rapid filtration: the water quality parameters of the influent water, marked pH, concentration of iron(II), homogeneous oxidation in the supernatant layer, surface sorption and heterogeneous oxidation kinetics in the filter, and adsorption characteristics of the filter media. Adsorption isotherm data collected from different Dutch remidation places show that Fe(II) adsorption may vary strongly between them, but generally increases with higher pH. The model has a sensitivity for (experimentally) determination of adsorption parameters and the heterogeneous oxidation rate.
Wang et al. [113] in astudy focused on Effects of solution chemistry on the removal reaction between Fe(II) and calcium carbonate-based materials by using a permeable reactive barrier consist of calcium carbonate-based materials (CCBMs), such as limestone. There is no significant effect on the uptake of Fe(II) by limestone from pH 7 to 9. Na+ significantly affected elemination of Fe(II) at levels of 100 mg/L and above. Ca2+ and Mn2+ showed effect on removal as low as 10 ppm Ca2+ and 5 ppm Mn2+. natural organic matter (NOM) premixed with Fe(II) (10 ppm Dissolved organic carbon (DOC) ) resulted in final Fe(II) levels above GCTL (groundwater cleanup target level). NOM retained 0.05 mg Fe(II)/mg for 2/3 sources and 0.032 mg/mg for 1/3.
Indah and Helard [114] conducted astudy on Evaluation of Iron and Manganese-coated Pumice from Sungai Pasak, West Sumatera, Indonesia for Fe (II) and Mn (II) Removal from aqueous model solutions. The effect of soaking time for iron and manganese coating was studied and as comparison. The experiments were performed in batch mode at room temperature between 20 and 25 oC, pH 7; adsorbent dose of 10 g/L; adsorbent diameters of 0.30-0.50 mm; 90 minutes of soaking time and100 rpm of agitation speed. The results showed that the optimum soaking time for manganese coating and iron for removal of Fe (II) and Mn (II) was 100 hours. Iron-coated pumice showed to have high removal efficiency compared to uncoated and manganese-coated pumice. More than 84% of Fe(II) with 15 ppm initial concentration was removed by 10 g/L iron-coated pumice, while by using uncoated and manganese-coated pumice, the elemination efficiencies were less than 75% . The desorption study noticed that up to 20% of Fe (II) was recovered from the three kinds of pumice adsorbent. Overall research indicated that pumice from Sungai Pasak may be a promising adsorbent for iron disposal from water and wastewater.
1.5 Enviromental pollution with lead metal and its removal:
Lead is a highly toxic metal whose widespread use has give rise to comprehensive environmental pollution and problems of health in many world parts. Lead is a bright silvery metal, slightly bluish in a dry atmosphere. It begins to tarnish with air contact, thereby forming a complex mixture of compounds, depending on the given conditions. Figure (7) shows various sources of lead pollution in the environment [115].
The lead exposure sources involve mainly industrial processes, smoking and food, drinking water and domestic water sources. The lead sources were gasoline and house paint, which has been extended to plumbing pipes, lead bullets, storage batteries, pewter pitchers, faucets and toys [116]. larger than 100 to 200,000 tons of lead per year is being emitted from car exhausts in the US. Some is taken up by plants, fixation to soil and flow into water bodies, hence human exposure of lead in the general population is either owing to drinking water or food. Lead is an extremely toxic heavy metal which disturbs different physiological processes of plant and unlike other metals, such as, copper, manganese and zinc, it does not play any biological functions. A plant with high concentration of lead fastens the reactive oxygen species (ROS) production, resulting in damage of lipid membrane that finally leads to destruction of photosynthetic and chlorophyll processes and suppresses the plant overall growth [117]. Some research stated that lead is capable of suppressing the tea plant growth by reducing biomass and debases the tea quality by changing the quality of its components [118]. Even at low concentrations, lead remediation was found to cause large instability in ion uptake by plants, which in turn leads to significant metabolic changes in photosynthetic capacity and ultimately in a strong inhibition of plant growth.
Figure (7): Various sources of lead pollution in the environment.
Poisoning of Lead was considered to be a classic disease and the marks that were seen in children and adults were fundamentally attached to the gastrointestinal tract and the central nervous system [119]. Lead poisoning can also take place from drinking water. The pipes which carry the water may be made of lead and its compounds which can pollute the water [120]. According to the Environmental Protection Agency (EPA), lead is regarded a carcinogen. Lead has large effects on different parts of the body. Distribution of Lead in the body initially based on the blood flow into various tissues and almost 95% of lead is precipitated in the form of insoluble phosphate in skeletal bones [121]. Toxicity of lead, also called lead poisoning, can be either chronic or acute. Acute exposure can result in headache, abdominal pain, appetite loss, renal dysfunction, vertigo, sleeplessness, arthritis, hallucinations, hypertension and fatigue. Acute exposure chiefly occurs in the work place and in some manufacturing industries which make use of lead. chronic exposure of lead can cause psychosis, autism, weight loss, allergies, mental retardation, dyslexia, hyperactivity, kidney damage, paralysis, muscular weakness, birth defects, brain damage, and may even lead to death [122].
Eventhough lead toxicity is preventable it still remains a dangerous disease which has effect on most of the organs. The plasma membrane moves into the brain interstitial spaces when the blood brain barrier is exposed to great levels of lead concentration, leading a condition called edema. It disrupts the intracellular second messenger systems and alters the the central nervous system functioning, whose protection is highly important. Domestic and environmental lead ions sources are the main reason of the disease but with appropriate precautionary measures it is possible to reduce the risk correlated with lead toxicity [120].
Generally, Impact of lead exposure in humans has been known to cause wide variety of health problems such as [123] :
• Various forms of blood disorders and Anemia
• Rapid deterioration of brain and the nervous system
• fertility decreasing both in men and women
• Failure of the kidney
• Alzheimer disease
Many studies have been reported for lead elemination from aqueous solutions. Pala and Dursun [124] studied that the results of a study on adsorption of Pb (II) ions from artificial contaminated tap water by using a natural zeolite (Clinoptilolite). Clinoptilolite mineral which has mesh size of 25-140 was used by activating with HCl, and the efficiencies of lead ion disposal were evaluated. Experiments were occured under laboratory batch conditions were run at different values of pH, temperatures. The highest efficiency of removal was found as about 87% at pH 5. In similar way, experiments were done at different temperature values, and the utmos efficiency was achieved at 30oC. The efficiency obtained under these conditions was 89.95%. The highest lead disposal efficiency was achieved with shaking speed of 200 rpm.
Mavropoulos et al. [125] in their study found that the composite of hydroxyapatite-alginate was effective in the elemination of lead ions and lead phosphate nanoparticles from high-polluted simulated gastric fluid. The cross-linked polymer chain had a double role: (i) keep Pb2+ ions and lead phosphate nanoparticles bounded to the surface of bead, impeding their bioavailability in stomach fluid; and (ii) delay dissolution of HA in the stomach acidic conditions, confirming that an excess of Ca2+ will not be released to simulated gastric fluid. Desorption studies in simulated enteric fluid stated that lead stayed immobilized in the calcium phosphate phase in the intestinal tract. These results indicate HA–alginate composite as effective system for heavy metals disposal from polluted gastric and enteric human fluids, reducing its adsorption by the human body.
Meski and his team [126] showed in their work with hydroxyapatite prepared from the egg Shell , that the carbonate hydroxyapatite prepared from egg shell (CHAPF) represents the highest capacity for Pb2+ ions adsorption from aqueous solution. It has been found that the initial adsorption rate was high. The sorption process obey the model of Langmuir isotherm with low temperature dependency and high adsorption capacities. The thermodynamic functions were calculated, and it can be concluded that the Pb2+ adsorption over CHAPF is an exothermic and spontaneous process. The adsorption was greatly pH dependent, with a high uptake of lead at pH = 3. These results show that the lead uptake by CHAPF was very sensitive to the initial concentration of Pb2+ in aqueous solution. For the high concentrations [(500 to 700) mg·L-1], two stages were observed: Pb2+ ions adsorption on the CHAPF surface and an ion exchange reaction between Ca2+ of CHAPF and Pb2+ ions in aqueous model solution.
Shrestha et al. [127] conducted a study on lead (II) disposal from aqueous solutions using prepared activated carbon. Two series of carbon have been synthesized from Lapsi seed stones by treating with concentrated H2SO4 and HNO3 in a mixture with H2SO4 in the ratio of 1:1 by weight for disposal of metal ions. pH 5 was the optimum pH for lead adsorption. For the equilibrium isotherms description, the adsorption data were better fitted with the Langmuir adsorption equations than Freundlich equation. The maximum adsorption capacity of Pb (II) on the produced activated carbons was 277.8 mg/g with a mixture of HNO3 and H2SO4 and 423.7 mg/g with H2SO4. The waste material used in activated carbons preparation is readily available and cheap. Therefore the carbons synthesized from Lapsi seed stones can work as potential low cost adsorbents for the elemination of Pb (II) from water.
Jalali [128] investigated astudy on stalk of Sunflower, an agricultural waste, acts as an adsorbent for the cadmium and lead disposal from aqueous solutions. Adsorbent was synthesized by washing residue of sunflower with deionized water until the solution become colorless. The results stated that the adsorbent has good sorption potential and maximum removal of metal was detected at pH 5. Within 150 min of operation about 97 of Pb ions were eleminated from the effluents. Curves of lead sorption were well fitted to the modified two-site Langmuir model. Lead adsorption capacities at optimum operation conditions were 182 mg/g. The kinetics of Pb ions adsorption from aqueous model solutions were also analyzed. The experimental data were foud to be fitted to pseudo-second-order kinetic model. fitted models (R2 > 0.999). The maximum adsorption capacity for Pb(II) ions adsorbed onto entrapped silica nanopowders was evaluated to be 83.33 mg/g.
Soltani and his team[129] conducted astudy on adsorption of Pb(II) ions from the aqueous solution by using entrapped silica nanopowders within calcium alginate in order that determination the thermodynamic, isotherm and kinetic of the adsorption process. According to the results, an initial pH of 5.0 was found to be optimal for the Pb(II) ions adsorption. The capacity of adsorption reached to 36.51 mg/g with increasing the contact time to 180 min at 50 ppm as initial Pb(II) ions concentration. However, the equilibrium contact was estimated to be 90 min owing to no significant increase in adsorption effeciency after this time. The results of studies stated that the isotherm of Langmuir and pseudo-second order model of kinetic were the best.
Yarkandi (2014) [130] carried out batch experiments for lead separation from waste water using natural american bentonite and activated carbon. The results show that the amount of Pb++ adsorption increases with solution pH, initial concentration of metal ion and contact time but decreases with temperatures and amount of adsorbent. The adsorption process has well fit pseudo-second order kinetic model. Langmuir and Freundich adsorption isotherm models were found to be applicable to the adsorption process where both were applies to analyze adsorption data. Thermodynamic parameters e.g. ΔH°, ΔG° and ΔS° of the adsorption process was found to be endothermic. Finally it can be seen that activated carbon was found to be less effective for disposal of Pb+2 ions than bentonite.
Bartczak et al [131] in a study investigating the lead (II) ions adsorption from aqueous model solution on peat as adsorbent with low-cost, observed that The sorption capacities of peat with respect to lead(II) ions was 82.31 mg(Pb2+)/g. Slightly well results were achieved with adsorption efficiency reached 100% just after 3 min (15 and 30 ppm), 5 min (50 ppm) or 15 min (100 ppm) of the process. It indicates the utmost affinity of the peat surface for lead ions. the optimum adsorbent mass was found to be 5 g/L. To prevent metal precipitation as hydroxides and including the obtained results pH = 5 confirmed as an ideal.
Sangeetha et al [132] investigated a study on lead ions disposal from aqueous solution by using novel hydroxyapatite/alginate/gelatin composites. Pb2+ elemination ability of wet precipitation synthesized biosorbents HA/Alg and HA/Alg/Gel has been investigated for different dosages. Complete disposal was obtained from 7 to 24 hours by both the adsorbents. The sorption kinetics were found to be best fit to the pseudo-second order equation and the equilibrium well followed Langmuir isotherm model. The elemination capacity was higher for lower dosage studied and the rate of disposal was higher for the higher dosage studied.The sorption mechanism involved was dissolution/precipitation, ion exchange and surface complexation process. according to The results, it can be seen that both the composites under study are potential candidates for Pb2+ disposal and precisely gelatin enhanced the maximum sorption than the alginate alone which is composed with the hydroxyapatite.
Cheraghi et al. [133] showed in their work with waste tea leaves, upon parameters optimization like initial metal concentration, Temperature, pH and adsorbent amount, that maximum uptake efficiency was achieved at pH 6. Also, as the initial metal concentration decreased the adsorption of Pb (II) ions increased. the equilibrium adsorption isotherm data fits well with the Langmuir isotherm model and its calculated maximum adsorption capacity of monolayer was 166.6 mg/g at 25±0.1˚C. The sorption kinetics were found to be best fit to the pseudo-second order equation.
Kanyal and Bhatt [134] reported the results of astudy on adsorption of Pb (II) ions from waste water by using household waste as an adsorbent. Banana peels, Pumpkins and Chicken eggshells are considered as good adsorbents for elemenation of heavy metals from polluted water. The effects of various parameters such as agitation speed, pH and residence time were examined and best results were observed at pH 7, 90 mins and 100 rpm. The results stated that household waste usage such as these can be act as a good biosorbent for disposal of heavy metals on a large scale and establish effective, and inexpensive techniques in wastewater remidation.
Wang et al. [135] carried out a study for selective disposal of lead and other metals such as cadmium and copper from wastewater by gelation with alginate for effective recovery of metal. The results evaluated that gels can be formed speedily between the metals and alginate in lower than 10 min and the rates of gelation fit well with the pseudo second-order kinetic model. The optimum ratio of dosing of alginate to the metal ions was found to be between 2:1 and 3:1 for Pb2+ elemination and around 4:1 for Cu2+ and Cd2+ elemination from wastewater, and the metal disposal efficiency by gelation increased with the solution pH. Alginate has a higher affinity of gelation toward Pb2+ than Cd2+ and Cu2+, which permitted a selective uptake of Pb2+ from the wastewater in the existing of Cd2+ and Cu2+ ions.
Liu et al. [136] focused on the disposal of lead ion from waste water using hydroxyapatite scaffolds synthesized from scales of fish. Powder of fish scale obtained from Tilapia fish (Oreochromis mossambicus) was used for preparing scaffolds for lead removal. The maximum adsorption capacities (qmax) were 344.8 mg/g and 208.3 mg/g in solutions pH of 2.2 and 5, respectively. More than 99.9% of the lead ion was eleminated after 20 min.
Al Lafi et al. [137] conducted astudy on Lead removal from aqueous aolutions by polyethylene waste/nano-manganese dioxide composite . The adsorption results investigated that the synthesized adsorbent can effectively eleminate Pb+2 ions from aqueous solutions, with a maximum adsorption capacity of 50.5 mg/g. Almost 60 % of the initial Pb+2 concentration were adsorbed within the first hour, and it was concluded that 2 h was the optimum time for Pb+2 elemination . A pH value of 5.0 was determined as an optimum and was used for the rest of this study. Regeneration of the composite can be performed using 0.5 mol/L HCl solution with Pb+2 percentage of recovery reached 95 %. It also efficiently adsorbed Pb2+ after five sorption/desorption cycles with 84 % as percentage of removal.
Heraldy [138] conducted a study of sorption of Pb(II) from the aqueous mediums using biosorbent synthesized from waste of tomato and residue of apple juice (AR). The optimum conditions for maximum removal percentage of Pb(II) by biosorbents were found to be 0.1 g of sorbent at pH 4.0 and 90 min contact time for tomato waste and 60 min for AR. The experimental data were found to be well matched with Freundlich than Langmuir isotherm model. A kinetic study showed that Pb(II) sorption follows the pseudo-second-order kinetics, which confirms that AR and waste of tomato biosorbents are 108 and 152 mg/g, respectively.
1.6 Role of alginate in heavy metals removal:
Alginate is belong to the anionic polymers family which is naturally occurring usually produced from brown seaweed, and has been widely investigated and used for many biomedical applications, due to its relatively low cost, low toxicity, biocompatibility, and mild gelation by divalent cations addition such as Ca2+ [139]. Alginates are a polysaccharides composed of variable ratios of β-D-mannuronate (M) and its C-5 epimer α-Lguluronate (G) linked by 1–4 glycosidic bonds (Fig.8). In the 1880s, alginates were first separated from brown seaweeds, and its production for commerce started in the early 20th century.
The production of alginate can be carried out by two bacteria genera, Azotobacter and Pseudomonas and various genera of brown seaweed and [140]. Alginate which is available in commerce can usually extracted from brown algae (Phaeophyceae), including Laminaria digitata, Macrocystis pyrifera, Laminaria japonica, Ascophyllum nodosum and Laminaria hyperborea by remidation with aqueous alkali solutions, usually with NaOH. The extract is filtered, and in order to precipitate alginate either calcium or sodium chloride is added to the filtrate. By treatment with dilute HCl, the alginate salt can be converted to alginic acid. After further purification and conversion, power of water-soluble sodium alginate is generated. [139, 141]
Figure (8): Chemical structure of alginate. M – mannuronate residues, G – guluronate residues.
The synthesised alginates were found to be extremely effective in uptake of heavy metal ions from aqueous model solutions. Removal of more than 93% Cr(VI) ions was obtained from aqueous solution in batch process using this type of biosorbent. [142, 143]
Alginate is regarded a biopolymer with many applications in food industry, drug delivery systems, cosmetics and cell encapsulation. In wastewater remediation could play a significant role in disposal heavy metal ions due its advantages, such as biodegradability, facile obtaining procedure, economical, biocompatibility and environmental friendly [144].
NiŃa et al. [144] in astudy found that calcium alginate microparticles has a good affinity for the divalent cations. The heavy metals adsorption was examined as a function of residence time between the samples of synthetic wastewater and the alginate and the polymeric beads morphology. The microparticles of calcium alginate were synthesized using a laboratory procedure, sodium alginate aqueous solution was dropped in a solution of calcium chloride. The polymeric microparticles which have a controlled porosity seem to be an appropiate alternative to develop a aprocess for elemination of heavy metal from industrial wastewater.
Singh et al. [145] investigated astudy on effective removal of Cu2+ ions from aqueous medium using alginate as biosorbent. Maximum removal of Cu2+ ions (85.3%) from aqueous medium was observed at pH 5.5, alginate dosage of 2.5% and initial copper concentration of 275 ppm with 50 min as agitation time. Thus, the reultant experimental data has been fitted well with both Langmuir and Freundlich isotherm models.
1.7. Heavy metals removal by bentonite and kaolin clays:
Clay is one of potential good adsorbent substitutes to other adsorbent owing to its layered structure, large surface area, mechanical and chemical stability and high capacity of cation exchange. The existence of two acidity types, Lewis and bronsted in clays increases the clays adsorption capacity. Aluminum oxides and clay minerals such as bentonite and kaolin, are the most wide-spread minerals of the earth crust which are known to be good adsorbent of different metal ions, organic ligands and inorganic anions [146, 147].
1.7.1. Bentonite:
Bentonite is considered a strong candidate as an adsorbent for heavy metal elimination due to its abundance and its low cost. Bentonite as a representative clay mineral is a clay chiefly consist of montmorillonite, a 2:1 type of aluminosilicate Bentonites are extremely valued for their sorption [146, 147].
Zia and his team [148], evaluated the removal of Pb2+ and Cd2+ by using Methionine modified bentonite/Alginate (Meth-bent/Alg) nano composite. The desorption study presents that 99% of the adsorbed Pb2+ and Cd2+ can be desorbed by using oxalic acid (0.1 M) as eluting agent with regeneration ability up to fifth cycle effectively.
Wu et al. [149] reported the adsorption of Th4+ from aqueous solution by using Novel magnetic organo-bentonite-Fe3O4 (polysodium acrylate) (OB-Fe3O4 PSA) superabsorbent nanocomposites. OB-Fe3O4 PSA super absorbent could be regenerated through the desorption of Th4+ using HCl solution (0.1 mol/L) and the adsorption capacity was still greater than 3.6 mmol/g after five successive adsorption–desorption processes.
Tan and Ting [150] reported that plain alginate and alginate immobilized bentonite beads have good reusability potential. remidation with HCl (10 mM) successfully eluted 93.05% and 94.33% of the Cu2+ ions loaded onto plain alginate and alginate immobilized bentonite, respectively, after three cycles of sorption–desorption test. There was no significant difference in the percentage of Cu2+ desorped in the three sorption–desorption cycles for both the plain alginate (93.15%, 93.54% and 92.48%) and immobilized-bentonite (92.38%, 96.03% and 94.75%). This established high reusability of the developed immobilized bentonite without remarkable losses in their Cu2+ disposal capacities.
1.7.2. Kaolin:
Kaolin is considered a type of clay rock, which incloses some chemical elements such as Na, Al, Mg, Ti, Ca, Fe, Ka, Si , and so on. The silicon mass proportion is more than 50%. Kaolin can be divided into two classes, which is coal-series kaolin and nature kaolinite The monolayer crystal structure of kaolin is comprised of siliconoxygen tetrahedral sheets and aluminum-oxygen octahedral sheets, just like the molecular structure model of kaolin shown in Figure (9) [151]. Metal ions removal using kaolinite clay is depend on mechanisms of ion exchange and adsorption and kaolinite has a relative low capacity of cation-exchange (CEC) [3–15 meq/100 g of clay] and smaller surface area ranged between 10 and 20 m2/g [152].
Figure (9): The molecular structure model of the kaolin.
Li et al. [153] reported that a novel environmental friendly material, calcium alginate immobilized kaolin (kaolin/CA), which synthesized using a sol-gel method, have good effeciency for copper uptake from waste water. the experimental adsorption was described using the Langmuir isotherm, the maximum capacity of Cu2+ adsorption by the kaolin/CA reached up to 53.63 mg/g. The thermodynamic studies indicated that the adsorption reaction was found to be an endothermic and spontaneous process.
Yavuz and his team [154] investigated the elemination of heavy metals such as Cu(II), Ni(II), Co(II) and Mn(II) from aqueous solution using raw kaolinite. The sorption of these metals on kaolinite conformed to Langmuir adsorption equation. Langmuir Cm constants for each metal were found as 0.919 mg/g (Co), 10.787 mg/g (Cu), 1.669 mg/g (Ni), 0.446 mg/g (Mn), at 25 oC, respectively. Also, kinetic and thermodynamic parameters like entropy (ΔS), enthalpy (ΔH) and free energy (ΔG) were evaluted and indicated that heavy metal adsorption on kaolinite was an endothermic process and the process of adsorption was preferable at elevated temperatures.
Larakeb (2017) [155] evaluted the Zinc Removal from Water by Adsorpion on Kaolin and Bentonite clays. The kinetics of adsorption results showed that zinc disposal is max. with and 45.48℅ efficiency for kaolin after 60 min of residence time and after 20 min with 89.8 ℅ efficiency for bentonite. Adsorbent dose increasing from 0.5 to 8 g/l enhance zinc elemination efficiency for 5 ppm like an initial concentration. Zinc disposal efficiency by the two adsorbent decreases with rising of the initial Zn concentration from 2 to 20 ppm. pH of treatment has considerable effects on the retention rate of zinc. The efficiencies of Zn removal are noticeable at basic pH. Whatsoever reaction parameter tested, it appears that kaolin is less effective than bentonite.
1.8. Removal of heavy metals by gelatin:
It has been reported that gelatin has agood affinity for heav metals removal, individaly or combined in composites with other materials. Itabashi et al. [156] found that The good effect of copper elimination by gelatin was achieved by the foam treatment of this solution. And also lead was successfully eliminated by the same treatment. gelatin powder was used for the adsorption treatment to raise this effect before the foam treatment. About 99 % of copper in the range of pH between 6.5 and 7.3 and approximately 100% of lead at pH 7.0 was eliminated respectively. Hayeeye et al [157] in their study found that, The maximum capacity of adsorption of gelatin/ activated carbon for Pb2+ ions was obtained to be 370.37 mg g-1. The separation process for Pb2+ ions was found to be relatively rapid with 92.15% of the adsorption finished in about 5 min as residence time in batch conditions. Adsorption was achieved at pH value as low as 2.0 and maximum adsorption was observed at a pH of almost 5.
1.9. Hydroxy apatite from bio waste materials:
Hydroxyapatite (HAP, Ca10(PO4)6(OH)2), a naturally available form of calcium phosphate and a component of hard tissues, has been reported to work as an efficient ion uptake material for different heavy metals from aqueous medium owing to its low solubility of water and excellent reactivity. The high stability of HAP structure, along with its flexibility permit a high variety of exchanges (particularly Ca ions with divalent heavy metal ions, such as As, Cd, Cu, Zn, Pb, Co, Ni, Sb,U, Hg, of huge importance in the environmental science field [158-160]. HAP can be extracted from different biowaste such as, eggshells and bovine bones. It can be synthesized through various methods which can be generally divided into two major routes: solid state reaction and wet methods [161]. including sol-gel technique, wet precipitation, hydrothermal process, mechanochemical method. Depending on the used techniques, HAP with several morphologies, composition, specific surface and crystalline degree have been obtained and appear to have different effects on the mechanical properties, bioactivity and dissolution behavior in biological environment [160, 162].
1.9.1. Preparation of hydroxyl apatite by sol-gel technique:
Sol-gel method is used for obtaining HAP powder of fine particle (nanoparticle size). In the sol-gel method of the HAP the calcium compounds and phosphorus precursors are transformed through condensation and hydrolysis reactions to the amorphous gels, which are further converted to ceramics when heated at comparatively low temperature. The polycondensation and hydrolysis are not separated in time, but occurs simultaneously [163]. Ceramic materials prepared by sol-gel route present many advantages over the others, such as homogeneous composition, low synthesis temperature and high product purity. Additional advantage of the sol-gel method is its applicability for surface coating.
There is no any form of secondary environmental damage as a result of high biocompatibility and its slightly- alkaline pH. The efficiency of HAP in eliminating heavy metal ions extremely depends on ion nature, diameter, charge and concentration, in addition to the treated water properties (temperature, pH) [160]
Putra and his team [164] showed in their work with eggshell, that for batch adsorption studies , at 90 min equilibrium time, 0.1 g biomass dosage and pH 6 were optimum biosorption conditions for Zinc and Copper ions elimination from aqueous mediums.
Agarwal and Gupta [165] in their study with eggshells focused mainly on evaluting varying concentrations (5, 10, 20, 40, 100 mg/L) of lead and copper ; this study reported a 92% - 100% removal of Cu when 0.5 to 1.5 g of eggshells (adsorbent) was used against 5 and 10 ppm of Cu; and adsorption efficiency of 80% to 100% for Pb at the same concentrations.
Deydier et al. [166] conducted a study for elemination of Pb from effluents using alow-cost material from meat and bone meal combustion residue. This residue was regarded as an apatite-rich material and was used as a low-cost substitute of hydroxyapatite in lead elemination from water . the mechanism was found to be as in pure apatite: surface complexation and dissolution of calcium hydroxyapatite, followed by lead hydroxyapatite precipitation.
Rohaizar et al. [167] in astudy, observed that pH = 7 and 350 rpm as an optimum agitation rate were ideal for copper elemination from water.
Avram et al. [160] in a study investigated that low crystallinity HAP which prepared by the direct reaction of diammonium hydrogen phosphate and calcium nitrate at alkaline pH, can be successfully used in heavy metal disposal from mine wastewater. For all the 10 metals studied (Zn, Cd, Pb, Mn, Co, Fe,Cr, Cu, Ni and Al), their content was fastly reduced by contact with HAP under the legal allowable limits for wastewater discharge in natural environment. The ion exchange importance in sorption processes was revealed and the pseudo-2nd order kinetics of manganese ions sorption on HAP was estimated.
1.10. Heavy metals removal by silica nanoparticles (Diatomeous):
Silica is used widely in nanoparticles coatings which used in water purification techniques. Silica coating activates the NPs surfaces having various functional groups owing to the abundant existing of silanol groups on the silica layer. It also prevents leaching low pH situations of NPs. It also facilitates the NPs with non- specific moieties, highly and group specific ligands. Polymer layered silicate nanocomposites possess improved properties at low filler contents. At neutral pH, as the particle size increase, the acidity of Si NPs will increase resulting in 5 to 20% ionisation of silanol groups, causing attraction between anionic Si surface and cations by ion pairing [168].
Surface-functionalized nanoporous silica, often referred to as self-assembled monolayers on mesoporous supports (SAMMS), has previously presented the ability to act as very effective sorbents for heavy metal uptake in a range of environmental and aquatic systems [169]. Diatomite is belong to the siliceous rock family, silicon dioxide is the main constituent of it with the proportions up to 90%. Diatomeous has some advantages such as wear resistance, heat resistance non-toxic and large specific surface area,..etc. Diatomite is a sort of polyporous material. The diatomite porosity is up to 90%, which means that diatomite has great adsorbability [151].
Soltani et al. [129] conducted astudy on adsorption of Pb(II) ions from the aqueous solution by using entrapped silica nanopowders within calcium alginate in order that determination the thermodynamic, isotherm and kinetic of the adsorption process. According to the results, an initial pH of 5.0 was found to be optimal for the Pb(II) ions adsorption. The capacity of adsorption reached to 36.51 mg/g with increasing the contact time to 180 min at 50 ppm as initial Pb(II) ions concentration. However, the equilibrium contact was estimated to be 90 min owing to no significant increase in adsorption effeciency after this time. The results of studies stated that the isotherm of Langmuir and pseudo-second order model of kinetic were the best.
Karnib and his team [170] used a composite from Activated Carbon, Silica and Silica Activated Carbon. Silica/AC (2:3) composite showed the greatest elemination percentage for 30 & 200 ppm nickel. SEM images revealed that AC was a microparticle with 25 μm as an average size, while silica were nanoparticles having an average size of 12 nm. Silica/AC (2:3) composite was the most effective microparticle for nickel disposal and it is highly recommended to be used in water treatment for its high adsorptive capacity followed by AC and silica nanoparticles.
Aim of the work
The presence of toxic heavy metals in water has caused several health problems with animals, plants, and human. So that the removal of toxic heavy metals from polluted waters are one of the most important issues of environmental remediation.
The development of new products which are abundant in nature, cheap and have no environmental impact for treatment of natural resources is an important area of material technology. Calcium alginate and its composites fulfills both characteristics and have the ability to eliminate heavy metals from industrial streams.
Hence the aim of the present work is to synthesize calcium alginate and different calcium alginate composites from clays and biowaste materials (egg shell and bovine bones) and their characterization using XRD, FTIR, EDX and SEM.
The second aim of this research is to use the prepared powder samples to remove two of most hazardous heavy metals (Pb2+ and Fe3+) and measure the efficiency of each sample for remediation process.
2. Experimental and Methods
2.1. Materials, Solutions and Chemicals:
Two different biowaste materials are used in preparation of Nano hydroxyapatite –calcium alginate composites are listed in Table (3).
Table (3): Raw biowaste materials and their sources.
Material Source
Egg shell Local Hen’s egg shells
Bovine bone Local Butcher shop (bovine femur bone)
High purity analytical grade chemical material have been used in the current study are listed in Table (4).
Table (4): Solutions and chemical materials used in the current study .
Materials Chemical composition Manufacture
Diammonium hydrogen phosphate (NH4)2HPO4 Oxford laboratory India
Ammonia solution NH4OH BDH Analar England
Lead nitrate Pb(NO3)2 BDH
Iron nitrate Fe(NO3)2.9H2O BDH
Calcium chloride CaCl2 ANALAR
NaOH solution NaOH WINLAP
HNO3 solution HNO3 WINLAP
Sodium alginate C6H7O6Na Oxford laboratory India
Gelatin - Oxford laboratory India
Raw material powders are prepared and their designation in the present study is given in Table (5).
Table (5): Designation of Raw materials source in the present study.
Raw material The Source
HAP (1) Egg shells calcined at 900 °C
HAP (2) Bovine bone calcined at 1000 °C
Bentonite Abu Zaabal Fertilizer & Chemicals Co.( originating from china )
Metakaoline Kaolin from Sinai Peninsula calcined at 800 °C
Diatomeous Kazakhstan
2.2 Preparation of calcium alginate composite powders:
2.2.1 Calcium alginate:
Calcium alginate powder is prepared by using controlled gellification method [171] with some modification reported by Daemi and Barikani [172]. CA Nanoparticles are obtained by addition of CaCl2 (0.05 M) to solution containing sodium alginate (3%) by mechanical stirrer at high stirring rates (Fig. 10). Six gram of polysaccharide is dissolved in 200 mL of deionized water with high rate stirring at room temperature for 1 h. After homogenization of sodium alginate solution by mechanical stirrer, the solution (1 litre) of 0.05 molar calcium chloride is added to the system. After 1 h of the rotation, it permitted to stand at room temperature for 24 hrs. prepared nanoparticles are purified by centrifugation for 30 min. The precipitate is washed and filtered three times using double distilled water to remove the adsorbed sodium and chloride ions. The filtered CA precipitate is dried at 60 °C for 12 hours in a dry oven. This dried solids is finely grinded and sieved below 63 μm before characterization and usage.
Figure (10): preparation of calcium alginate nanoparticles
2.2.2. Bentonite:
Bentonite of Abu Zaabal Fertilizer & Chemicals Company which originating from china (Table 5) is used as a raw material in the preparation of CA – Bentonite composite. Natural bentonite dried in the oven with a temperature of 80oC for 12 hrs with the aim to eliminate moisture, bentonite ground with mortal to break chunks of bentonite then calcinied at a temperature of 800oC for 2 hours which aims to eliminate Cl bond on bentonite, and the results in the form of nanoparticles of bentonite [173].
2.2.3. MetaKaolin:
Kaolin from Sinai Peninsula is used as a raw material in the current study. Kaolin is calcined at 800 oC for 2 hrs to obtain Metakaolin (Table 5) that used in preparation of CA – Metakaolin composite.
2.2.4. Diatomeous silica:
Diatomeous which originating from Kazakhastan (Table 5) is used as a raw material in the preparation of CA – Diatomeous composite. Diatomeous is dried at temperature of 80 oC for 12 hrs with the aim to eliminate moisture.
2.2.5. Hydroxy apatite:
HAP (I)
The Egg shells mainly contain calcium carbonate (91% - 94%), calcium phosphate (1%) and other organic matters, which makes it preferable for synthesizing CaO [174]. The Nano-hydroxyapatite HAP (I) is prepared from hen’s egg shells by a method described by Laonapakul [175] with little modifications. About 50 gm. of egg shells is boiled for 30 minutes in hot water and the protein membrane is removed manually. Egg shells are dried at 80 ºC for 6 hrs. Dried eggshells solid are grinded in the agate mortar into a fine powder. The fine eggshells powder is calcined at 900 ºC for two hours in order to remove any organic residue. At this temperature the eggshells convert into calcium oxide (CaO), according to the following reaction:
Ca CO3 + Heat → CaO + CO2↑
Calculating the stoichiometric of Ca / P molar ratio = 1: 0.67 solution is prepared from CaO and DAP. Firstly, about 8.4 gram of CaO is dissolved in 100 ml 2M HNO3 and then 11.885 gm of DAP is added dropwise to calcium solution while stirring and maintaining a stoichiometry of Ca/P ratio of 1.67. NH4OH solution is added dropwise to the mixture. The pH of the solution is maintained at pH=10. A white precipitate solution is obtained and vigorously stirred for 30 minutes and permitted to stand at room temperature for 24 hours. The left white precipitate is filtered and washed three times using double distilled water to remove the adsorbed ammonia and nitrate ions. In order to obtain the final HAP solid, the filtered hydroxyapatite precipitate cake is dried at 80 °C for 10 hours in a dry oven. This dried powder is heated at 700 °C for 2 hrs. in air using control electric muffle furnace, employing a heating rate of 10 °C/min. Afterwards, the calcined powder is finely grinded and sieved below 63 μm before characterization and usage.
HAP (II)
One of the raw materials in the present study is bovine femur bones obtained from local butcher market. HAP obtained from bovine bones is prepared by a method conducted by Agnieszka et al. [176] with little modifications .In the beginning, the bovine bones (about 150 gm) are crushed into small pieces (1-2 cm) and then boiled in 1 M NaOH for one hour then in hot water for 1.5 hrs. for defatting and easier removal of the organic residues and macroscopic adhering impurities. The bones are washed and cleaned well with water and for several times afterwards. The process is followed by drying the bones at 80°C for 6 hrs. to evaporate the adsorbed water. The solid bone pieces is calcined at 1000°C for 2 hrs. at heating rate 10°C /min. and afterwards are cooled slowly to room temperature. The final solid is grinded and sieved below 63μm, and kept for characterization and usage.
2.2.6. Calcium alginate - Composites (CACS):
The compsites are prepared by mixing the CA with approprate amounts of additaves (Bentonite, Metakaolin, Gelatin, HAP(1), HAP(2) and Diatomeous) with 2:1 ratio, respectively. 6 gram of polysaccharide (sodium alginate) is dissolved in 200 mL of deionized water with high rate stirring at room temperature for 1 hrs. After homogenization of sodium alginate solution by mechanical stirrer, three grams of the chosen additives is add, then 1 litre of 0.05M calcium chloride solution is added. After 1 hour of the rotation, it permitted to stand at room temperature for 24 hrs. The prepared nanoparticles are purified by centrifugation for 30 min. The precipitate is washed and filtered three times using double distilled water to remove the adsorbed sodium and chloride ions. The CA- composites precipitates are dried at 80 °C for 24 hours in a dry oven. This dried solid are finely grinded and sieved below 63 μm before characterization and usage. Table (6) shows The synthesized composites and its abbreviations.
Table (6): The synthesized composite:
Composite Compound Abbreviations
1 Calcium alginate-Bentonite CAB
2 Calcium alginate-Metakaolin CAMK
3 Calcium alginate-HAP(1) CAHA(I)
4 Calcium alginate-HAP(2) CAHA(II)
5 Calcium alginate-Diatomeous CAD
6 Calcium alginate-Gelatin CAG
2.3. Chemical analysis of raw materials:
A- Diatomeous Silica:
Table (7): Chemical composition of Diatomeous used in the present study (XRF fused bed)
Compound Wt. %
SiO2 71.50
Al2O3 10.40
CaO 0.74
Fe2O3 3.66
MgO 1.23
SO3-- 0.71
Na2O 0.78
K2O 1.25
Cl- 0.28
TiO2 1.04
P2O5 0.10
Mn2O3 0.02
Total 91.71
Loss on ignition 8.10
Total 99.80
B- Bentonite:
Table 8: Chemical composition of bentonite used in the present study (XRF fused bed)
Compound Wt. %
SiO2 55.11
Al2O3 17.27
CaO 0.99
Fe2O3 9.03
MgO 2.27
SO3-- 0.34
Na2O 3.67
K2O 1.19
Cl- 0.62
Total 90.49
Loss on ignition 9.42
Total 99.91
C- kaolin:
Table (9): Chemical composition of kaolin used in the present study (XRF fused bed)
Compound Wt. %
SiO2 47
Al2O3 37
CaO 0.20
Fe2O3 0.20
MgO 0.02
Na2O 0.15
K2O 0.04
TiO2 1.30
Total 85.91
Loss on ignition 13.40
Total 99.31
2.4. Preparation of heavy metal ion solutions:
The metal cations of lead (Pb2+) and Iron (Fe3+) are used in the present study in the form of salts: Pb(NO3)2 and Fe(NO3)3.9H2O. One liter stock solution of each metal cation is prepared using double distilled water with metal ion concentrations of 100 ppm.
2.5. Heavy metals uptake reaction:
The metal cations reactions are conducted as follow: 20 mg of each calcium alginate composite solids is equilibrated for different time periods (5, 10, 15, 20, 30, 60 and 120 min. respectively) in glass vials with 10 ml metal cation solution with continuous shaking. After different time intervals, the solid phases are separated by centrifugation, the supernatant solution was collected for chemical analysis using Perkin Elmer 2380 Atomic Absorption Spectro Photometer (Figure 11).
Fig. 11: Perkin Elmer 2380 Atomic Absorption SpectroPhotometer.
2.6. characterization of Composite solids:
The prepared Composite solids are characterized by using Scanning Electron Microscope (SEM), X-ray diffraction (XRD), Fourier Transform Infrared (FTIR) spectroscopy and Energy Dispersive Analysis X-ray (EDX) techniques.
2.6.1. X-ray diffraction (XRD):
An X-ray diffractometer (Philips X’ PERTMPD, America, with Cu Kα radiation, 40KV and 30mA) is used to determine the mineral phases and crystallinity of the different composite powders; (Figure12).
Fig. 12: Philips X-Ray Diffractometer
The specification criteria of XRD are adjusted at 2Ө range = 5° - 60° and λ = 1.54 Ǻ at a scanning speed of 2° /min. Each composite is used to fill the aluminum mold of the diffractometer with an average thickness of about 10 mm. The obtained phases are identified by correlation with the corresponding joint committee on powder diffraction standard card (JCPDS). The average crystallite size (D) of the obtained composite powders is calculated from XRD using the Scherrer formula [177] as shown below.
where:
λ= the wave length of the X-ray.
β = the full width at half maximum (FWHM) of the peak at the maximum intensity
Ө = the diffraction angle
2.6.2. Attenuated Total Reflection Fourier Transform Infrared (ATR-FTIR):
(ATR-FTIR) technique is used to determine the main constituent chemical functional groups of the different prepared samples and the type of chemical bonding between the different atoms existing in the groups. ATR-FTIR spectrometer (Bruker, Germany Alpha-p) is configured with ATR-FTIR sample cell including a diamond crystal with a scanning depth up to 2μm. Sample powders are applied to the surface of the crystal then locked in placed with a”clutch – type” lever before measuring. The excitation of the corresponding elections when subjected to IR radiation is reflected in the spectrum as absorption bands at wavelength range from 4000 - 400 cm-1 at scanning speed of 2 cm-1 (Figure 13).
Fig. 13: ATR-Fourier Transform Infrared Spectrophotometer
2.6.3. Scanning Electron Microscope (SEM) and Energy Dispersive Analysis
X-ray (EDX):
Scanning Electron Microscopy (SEM) Model Quanta 250 FEG (Field Emission Gun) attached with EDX Unit (Energy Dispersive X-ray Analyses), with accelerating voltage 30 KV, a magnification of 14x up to 1000000, Gun.1n. FEI Company, Netherlands (Figure 14) is used to examine the surface morphology of the different prepared composite powders. The investigated samples are coated with gold (conductive layer) before imaging using EMITECH K550k sputter coater England. EDX analyzer is used to detect the chemical composition of the synthesized powders. The EDX system has a super ultra-thin window which means that it can analyze a wide range of elements.
Fig. 14: Scanning Electron Microscope with EDX
2.7. Adsorption studies
An accurately 0.02 g of CACs is added into 10 mL of solutions in a 25.0 mL glass tubes containing the specified concentrations of metal ions. The mixtures are shaken at 25.0 oC for a fixed period (2 hrs.) and at the end of shaking periods, the contents are filtered through filter paper. The filtrate is analyzed for final metal concentration using Perkin Elmer 2380 Atomic Absorption Spectro Photometer. Each experiment is performed in triplicate and the average of the results is recorded.
The effects of contact time (5-120 min.) and metal ion concentration on the sorption process are realized using the same methodology. The amounts of metal ion adsorbed onto CACs sorbent, qe (mg g–1), are calculated using the Equation:
qe = (Co – Ce) V/ m
where Co and Ce (mg L–1) are initial and equilibrium concentrations of metal ions, m (g) is the weight of sorbent in the solution and V (L) is the volume of the solution.
The efficiency of adsorption (removal %) is calculated according to the Equation:
% Removal = [(Co – Ce) /Co] × 100
2.7.1. Effect of contact time:
To study the effect of contact time on the adsorption efficiency of Pb2+ and Fe3+ ions by CACs, 0.02 g of CACs is added into 10 mL of 100 mg L-1 Pb2+ (pH 5.7 and 4) and Fe3+ ( pH 2.6) solutions at 25±1°C with time interval from 5 -120 minutes.
2.7.2. Effect of pH:
The effect of pH on the adsorption is performed only for Pb2+. The study is achieved with two pHs values (4 and 5.7) the original solution of lead is at pH 5.7 and to get pH 4 we use HNO3 solution (0.01 M). This process is conducted at 30 minutes contact time, 20 mg dosage of the different composites and 100 ppm of metal solution at 25±1°C. With respect to Fe3+ solution, the experimental work is carried out at pH =2.7 only. This is due to the precipitation of Fe3+ at pH ≥ 3.
2.7.3. Effect of dosage:
To study the effect of dosage (10, 15, 20 mg) on the adsorption efficiency of Pb2+ on CA-Np (as a standard model). The previous dosages are added to 10 mL of 100 mg L-1 Pb2+ at 25±1°C with 30 minutes contact time.
2.8. Adsorption kinetics:
Adsorption kinetic studies are important since they describe the solute uptake rate which controls the residence time of adsorbate at the solid–liquid interface and also provide valuable insights into the reaction pathways.
Pseudo–first–order (Equation I) and pseudo–second–order (Equation II) models were applied in order to investigate the adsorption kinetics of Pb2+ and Fe3+ ions onto CACs. The conformity between experimental data and the model-predicted values is expressed by the correlation coefficients (R2). Meanwhile, the capacity values calculated from the pseudo–first and second–order models are compared with that obtained from the experimental data. The kinetic models can be presented as follows,
ln (qe –qt ) = ln qe – k1 t (I)
t/qt = 1/k2qe2 + t/qe (II)
where qt is the amount of metal ion adsorbed (mg g-1) at time (t), qe is the maximum adsorption capacity (at equilibrium) (mg g-1) and k1(min−1) is the rate constant for pseudo–first–order sorption, qe is the maximum adsorption capacity (at equilibrium) (mg g-1) and k2 (g mg-1 min-1) is the rate constant for the second-order sorption. from plots of t/qt versus t for the second–order reactions, the k2 and qe values are calculated by using the values of intercept and slope.
2.9. Adsorption isotherm models:
The adsorption equilibrium is investigated by using two well-known isotherm models (Langmuir and Freundlich) to provide the fundamental physicochemical data and investigate the applicability of sorption process at a fixed temperature. The equilibrium conditions of the adsorption process are described by utilizing the linearized equations indicated below,
Ce/qe =Ce/qm + 1/ KL qm
where qm (mg g–1) and KL (L mg–1) are constants in Langmuir’s equation which are referred to the maximum adsorption capacity and the Langmuir model constant that is indirectly related to the energy of adsorption. Also qe and Ce parameters represent the equilibrium adsorption capacity and the equilibrium concentration heavy metal ion, respectively.
Freundlich adsorption isotherms assumes aheterogeneous surface with anon-uniform distribution of heat of adsorption over the surface with the possibility of the multilayer adsorption [178, 179].the Freundlich equation which is expressed as
qe =log KFCe1/n
where KF is the measure of adsorption capacity and n is the adsorption intensity linear form of Freundlich
log qe =log KF + (1/n) log Ce
where qe is the amount adsorbed (mg/g), Ce is the equilibrium concentration of adsorbate (mg/L), and KF and n are the Freundlich constants related to the adsorption capacity and adsorption intensity, respectively. A plot of log qe vs. log gives a linear trace with a slope of 1/n and intercept of log KF.
3. Results and Discussion
3.1. characterization of the synthesized composites.
3.1.1. Fourier Transform Infrared (FTIR) analysis:
The Fourier transform infrared (FTIR) spectra of calcium alginate and calcium alginate composites are given in Figures (15-18). Also their frequencies of the absorbtion bands of FTIR spectra and their structure assignments are given in Table (10). Spectrum of calcium alginate (Figure 15), showed important absorption bands regarding hydroxyl, ether and carboxylic functional groups. Stretching vibrations of O–H bonds of alginate appeared in the range of 3000–3600 cm-1 particualy at 3444 cm-1. Stretching vibrations of aliphatic C–H are observed at 2920–2850 cm-1. The observed bands at 1632 and 1454 cm-1 are attributed to asymmetric and symmetric stretching vibrations of carboxylate salt ion, respectively. 1154 cm-1 (CO-stretching of ether group) and 1025 cm-1 (C-O stretching of alcohol group)[180].
The spectra of CAG (Fig.16), the absorption band at around 3442 cm-1 concerned with OH stretching vibration for CA slightly broadened and shifted to a lower wave number with the blending with gelatin, suggesting the formation of an intermolecular hydrogen bond [181]. The strong absorption band at 1631 cm-1 for CA assigned to the asymmetric stretching vibration of COO- has coupled with the absorption band at 1631 cm-1 in gelatin.
CA raw material, reveals asimilar spectra with CAB as shown in Fig 17, in addition some peaks are found at 3699 and 3444 cm-1 are due to lattice OH and bound water stretching vibrations. A strong and sharp band is detected at 1022 cm-1 which is related to Si–OH stretching vibrations [182]. Peaks found at 1384 cm-1 is due to CO3 stretching of calcite, 1034 cm-1 assigned to Si-O stretching, and 875 cm-1 is due to OH bending of the Al-Al-OH group. A similar OH bending vibration is observed for Al-Mg-OH at 842 cm-1, 690 cm-1 assigned to quartz. Also, there is a shoulder peak at 520 cm-1 (Al-O-Si bending), and 464 cm-1 (Si-O-Si bending)[183].
The IR spectra of CAMK and CAD composites revealed a similar spectra with CAB composite spectra (Figure 17). CAD showed a strong band at 1084 and 1048 cm-1 due to Si–OH and Si-O vibrations. These bands overlabed with the band at 1025 cm-1 due to CO-stretching of alcohol group in alginate. A strong and sharp peak at 470 cm-1 are also detected due to Si-O-Si bending. The band of Si-O-Si bending in CAD is the strongest and sharpest compared to CAMK and CAB.
The HAP-Alginate samples (CAHA(I) and CAHA(II)) (Figure 18) revealed a similar spectra, at 1625-1630 cm-1 and 3440-3445 cm-1 (due to the presence of free water), 1454 and 874 cm-1 ( due to CO32- ions), 3570 and 630-633 cm-1 (due to structural OH of hydroxy apatite). These peaks due to the hydroxyapatite phase [184, 185]. The most intensive bands in the range of 1044 –1090 cm-1 corresponded to the triply degenerated asymmetric stretching vibrations of P-O. Otherwise the peak at 962.97 cm-1 indicates the non-degenerated asymmetric mode of PO43-. The very strong and sharp bands observed at 569-572 and 602 -603 cm-1 attributed to triply degenerated bending mode of the O-P-O in PO43- group. The larges parting distance of these bands revels the crystalline phase [184, 185].
Table (10): Assignments of the absorption bands of the IR spectra λ (cm-1) of the prepared composites
Peak Assignment Strength CA CAB CAMK CAHA(I) CAHA(II) CAD CAG
Structural OH of addittives w - 3699 - 3571 3570 - -
OH stretching mode of adsorbed water molec. or OH of alginate lattice structure s , b 3444 3444 3444 3444 3444 3443 3442
Stretching mode of aliphatic C–H. w 2924 2923 2924 2923 2924 2924 2924
asymmetric stretching vibrations of carboxylate salt ion m 1631 1634 1631 1629 1630 1629 1631
symmetric stretching vibrations of carboxylate salt ion vw 1427 1461 1431 1433 1419 1433 1427
CO-stretching of ether group vw 1155 1150 1153 - - 1103 1155
CO-stretching of alcohol group w 1024 1033 - - - - 1024
Si–OH stretching vibrations mode w - 1080 1081 - - 1084
CO3 stretching of calcite w 1384 1384 1384 1384 1384 1384 1384
Si-O stretching w - 1033 1052 - - 1048 -
OH bending of the Al-Al-OH group vw - 875 850 - - 845 -
OH bending vibration of Al-Mg-OH group of quartz vw - 842
690 777 - - 797 -
Al-O-Si bending vw - 520 510 - - 526 -
bending vibration mode of Si-O-Si s - 464 456 - - 470 -
vibration mode CO32- vw - - - 1454
874 1457
874 - -
structural OH in HA w - - - 631 631 - -
triply degenerated asymmetric stretching vibrations of P-O vs - - - 1044
1090 1048
1089 - -
non-degenerated asymmetric of PO43- vw - - - 962 962 - -
triply degenerated bending mode of PO43- ms - - - 602
569 602
571 - -
s = strong w = weak vw = very weak b = broad m = meadium
Figure 15: FTIR Spectra of Calcium alginate.
Figure 16: FTIR Spectra a- CA, b- CAG
Figure 17: FTIR Spectra a- CA, b- CAB, c- CAMK ,d- CAD.
Figure 18: FTIR Spectra a- CAHA(I) , b-CAHA(II)
3.1.2. X-ray diffraction for crystal phase detection:
Figure(19) presents the X-ray diffraction pattern of CA and CA composites in the range of 2θ = 5-60o diffraction degree . Two typical peaks in 2θ =16° and 22° are observed for calcium alginate. The XRD of CAG shows typical peaks around 12° and 21°[186].
The samples of CAB, CAMK and CAD powders revealed very similar XRD pattern peaks of quartz phase. These composites showed the maximum relative intensity (I/I0) peak of 100% quartz in the rang of 2θ =26.65o and 20.85o and d (Ao) value spacing equal 3.34 and 2.8, (reference code :01-070-3755). On the other hand, additional peaks are detected due to the presence of minor amounts of calcite (CaCO3) and halite (NaCl). The presence of calcite phase may be attributed to carbonation of calcium during composite synthesis, while the presence of NaCl may be due to the reaction of Na+ ions of alginate with Cl- ions of CaCl2 solution and its trapping bettween the layers, which can not completely removed by washing.
The well resolved XRD peaks of The sample HAP(I) ,CAHA(I), HAP(II) and CAHA(II) could be easily indexed on the basis of hexagonal crystal system of space group P63/m with respect to JCPDS file no. 9-432. They also revealed very similar XRD pattern peaks of 100 % HAP (reference code :01-086-1194) in the rang of 2θ =31o and 32o and d (Ao) value spacing equal 2.81 and 2.78. There is no any considerable shifts in 2θ are detected between HAP and HAP composite . The diffraction peaks of HA and CA-HAP exhibit sharp diffraction peaks which indicate the high crystallinity of the structure and there is no any additional phases are detected.
XRD analysis of HAP(II) and CAHA(II) also indicated the absence of secondary phases, such as tri calcium phosphate (TCP) or calcium oxide (CaO). In the case of HAP(I) and/or CAHA(I), their diffraction patterns revealed additional phase of β-tricalcium phosphate, beside the HAP as a main phase.
Hydoxyapatite prepared from bovine bone contain certain amount of carbonate (CO32-) in its lattice structure [175]. So the sample CAHA(I) are expected to be a carbonated apatite type, and carbonate ions affect on the the degree of crystallinite. For this reason, CAHA(II) ( HA prepared from bovine bones) revealed higher crystal size (83 nm) than CAHA(I) which prepared from eggshells (59 nm). There is no significant differences in the crystal sizes of the prepared HAP and HAP Compsites. Table (11) showed the crystal size, 2θ, d-spacing(oA), maximum relative intensity (I/Io) peak and the main phase detected for the prepared adsorbents
Table (11): 2θ, d-spacing(oA), Crystal size ( nm ) of the maximum relative intensity (I/Io) peak and the main phase detected for the prepared adsorbents.
Character
2θ
d-spacing (Ao)
Crystal size (nm)
Main phase detected
CAHA(I)
31.83
2.81
59
HA
CAHA(II)
31.77
2.81
83
HA
CAB
26.65
3.34
83
Quartz
CAMK
26.68
3.34
84
Quartz
CAD
26.63
3.34
59
Quartz
Figure 19: X-ray diffraction pattern of the prepared calcium alginate composite powders.
Figure 19: continue
Figure 19: continue
Figure 19: continue
3.1.3. Scanning Electron Microscope (SEM) and Energy Dispersive X-ray (EDX)
The SEM images and EDX elemental analysis of synthesized calcium alginate and calcium alginate composites are shown in Figure (20). The images are taken at 300× and 2500× magnification. SEM investigations showed the presence of agglomerates of irregular shape particles.
According to the results obtained by EDX analysis, carbon, oxygen and calcium are the major constituents of all the prepared composites related to calcium alginate polymer present, Also, Almost prepared composites confirmed the presence of calcite and halite as additional phases in minor amounts which is in aggreement with X-ray results.
EDX analyses of CA and CAG revealed that the major elements present are carbon and oxygen corresponding to polymer composition. The SEM of CAG (Fig.20(i)) showed a smooth and homogeneous morphology, suggesting high miscibility and blend homogeneity between calcium alginate and gelatin.
In the case of CAB, CAMK and CAD, the presence of an obvious peak related to the Si compounds is evident. According to the results obtained by EDX analysis, weigh percent of elements present indicating a large portion of the composites is composed of Si compounds which is suitable for an efficient sequestering metal cations from aqueous solution[129].
EDX analyses of the prepared HAP revealed that inorganic phases of bovin bone and egg shells were mainly composed of calcium and phosphorus as the major constituents with some minor components such as C, O, Na, Mg and Si. The weight and atomic percentage shows that the Ca/P ratio around 1.7 and 1.8 which is below 2 and acceptable where the ideal Ca/P ratio of HA is 1.67[174], in HAP composites (CAHA(I) and CAHA(II)) this ratio increased than 2 due to the excess amount of Ca crosslinkage of calcim alginate presents.
a) CA
Element Wt % At %
C K 35.08 46.27
O K 45.62 45.17
Na K 2.08 1.43
Cl K 6.18 2.76
Ca K 11.05 4.37
Total 100 100
b) CAB
Element Wt % At %
C 16.29 25.90
O 37.59 44.87
Na 1.14 0.95
Al 13.61 9.77
Si 16.26 11.06
Cl 6.83 3.68
Ca 7.08 3.37
Ti 1 0.4
total 100 100
Figure 20: SEM images and EDX analysis of prepared samples a) CA & b) CAB .
c) CAMK
Element Wt % At %
C K 9.75 16.10
O K 41.60 51.55
Na K 1.07 0.93
Al K 4.43 3.26
Si K 26.61 20.91
Cl K 6.25 3.50
Ca K 5.85 2.90
K K 0.98 0.49
Mg 0.45 0.37
Total 100 100
d) CAD
Element Wt % At %
CK 27 39.27
OK 39.73 43.39
Na K 2.46 1.87
Al K 2.30 1.49
Si K 5.16 3.21
Cl K 7.79 3.84
Ca K 14.62 0.21
K K 0.46 0.21
Mg 0.47 0.34
Total 100 100
Figure 20: SEM images and EDX analysis of prepared samples c) CAMK & d) CAD.
e) HAP(I)
Element Wt % At %
C K 4.66 9.69
O K 29.05 45.31
P K 20.45 16.47
Ca K 45.84 28.53
Total 100 100
f) HAP(II)
Element Wt % At %
C K 3.96 8.22
O K 28.88 45.05
Na K 2.42 2.62
Mg K 0.76 0.78
P K 18.95 15.27
Ca K 45.03 28.04
Total 100 100
Figure 20: SEM images and EDX analysis of prepared samples e) HAP(1) & f) HAP(2).
g) CAHA(I)
Element Wt % At %
C K 20.69 33.23
O K 36.42 43.90
Na K 1.74 1.46
P K 10.18 6.34
Cl K 2.62 1.42
Ca K 28.35 13.64
Total 100 100
h) CAHA(II)
Element Wt % At %
C K 7.32 14.25
O K 31.95 46.71
Na K 0.70 0.71
Mg k 0.40 0.39
P K 17.06 12.86
Cl K 2.71 1.79
Ca K 39.86 23.27
Total 100 100
Figure 20: SEM images and EDX analysis of prepared samples. g) CAHA(I) & h) CAHA(II).
i) CAG
Element Wt % At %
C K 39.21 49.41
O K 44342 44.07
Na K 2.08 1.43
Cl K 3.15 1.8
Ca K 11.05 4.37
Total 100 100
Figure 20: SEM images and EDX analysis of prepared samples. i) CAG
3.2. characterization of the composites after metal ion uptake:
3.2.1. FTIR Analysis after metal ion uptake:
FTIR spectra of CA and CAHA(I) after targeted metal ions uptake ( Pb2+, Fe3+ ) are shown in Figures 21 and 22 respectively and their peak assignments are represented in Table (12). The result of FTIR spectra showed that there is no new absorption peaks were detected . There are alittle peak shifts which may be attributed to the corporation and substitution of metal ion in the lattice structure of CA and CAHA(I).
Table 12: Change in the absorption bands of the IR spectra λ (cm-1) of CA and CAHA(I) powder after metal ion uptake
Peak Assignment Strength CA
CA
+
Pb2+
CA
+
Fe3+ CAHA(I)
CAHA(I)
+
Pb2+
CAHA(I)
+
Fe3+
Lattice Structural OH of hydroxyapatite w -
- - 3571 3571 3571
OH stretching mode of adsorbed water molec. or OH of alginate lattice structure s , b 3444 3444 3434 3444 3444 3444
Stretching mode of aliphatic C–H . w 2924 2924 2923 2923 2924 2923
asymmetric stretching vibrations of m 1631 1598 1631 1629 1599 1615
symmetric stretching of COO- ion vw 1427 1425 1424 1433 1433 1433
CO-stretching of ether group. vw 1155 1156 1155 -
CO-stretching of alcohol group w 1024 1021 1021 -
CO3 stretching of calcite w 1384 1384 1383 1384 1384 1381
vibration mode CO32- vw - 1454
874 1454
874 1458
874
structural OH in HA w - 631 631 631
triply degenerated asymmetric stretching vibrations of P-O vs - 1044
1090 1045
1090 1045
1090
non-degenerated asymmetric of PO43- vw - 962 962 962
triply degenerated bending mode of PO43- ms - 602
569 602
570 601
568
s = strong w = weak vw = very weak b = broad m = meadium
Figure 21: FTIR Spectra a- CA , b-CA+ Pb2+ , c- CA + Fe3+
Figure 22: FTIR Spectra of a- CAHA(I) , b- CAHA(I)+ Pb2+ , c- CAHA(I)+ Fe3+
3.2.2. X-ray diffraction for crystal phase detection after metal ion uptake:
XRD patterns of CA and CAHA(I) after Pb2+ and Fe3+ metal ions removal did not revealed any new phases. supported the proposal that Pb2+ and Fe3+ ions uptake was not dependent on dissolution/precipitation mechanisms. Pb2+ and Fe3+ ions removal may be occurs by adsorption mechanisms like surface complexation or ionic exchange [187]. As it can be seen in Figures 23-24, XRD patterns showed some changes in their relative intensities and crystal sizes (Table 13). Also, ther are some little shifts in d- spacing values, this may be due to the ion exchange between Ca2+ and metal cations of Pb2+ and Fe3+ in lattice structure of CA and CAHA(I) nanopowders.
Table 13: 2θ, d-spacing(Ao), Crystal size ( nm ) of the maximum relative intensity (I/Io) peak of CAHA(I) before and after metal ion uptake.
Character
2θ d-spacing (Ao) Crystal size (nm)
CAHA(I)
31.83 2.811 59
CAHA(I) + Pb2+
31.74 2.819
60
CAHA(I) + Fe3+ 31.71 2.821 37
Figure 23: X-ray diffraction pattern of CA after metal ion uptake
Figure 24: X-ray diffraction pattern of CAHA(I) after metal ion uptake
3.2.3. Scanning Electron Microscope (SEM) and Energy Dispersive X-ray (EDX) after metal ion uptake:
SEM of CA and CAHA(I) after metal ion uptake revealed some changes in morphology and microstructure(Fig. 26,27). Morever EDX results indicated the presence of Pb2+ and Fe3+ ions with CA and CAHA(I). As can seen in Figures 26 and 27, the peak of Fe3+ is more intense than that of Pb2+ and this confirms, the higher value of Fe3+ ions uptake by CA and CAHA(I) comparing with Pb2+ ions. The weight and atomic percentage of Ca2+ ions in CA and CAHA(I) after metal ion uptake was less than its value before metal ion removal as listed in Table (14). The decrease in calcium percentage may be attributed to ion exchange between targeted metal ions and; (1)calcium ions of calcium alginate in CA and CAHA(I) as follows:
Ca(ALG)2 + Pb2+ Pb(ALG)2 + Ca2+
(2) or Ca ions of HA present in CAHA(I). This ion exchange mechanism between Pb2+ ions (as example) and Ca2+ ions of HA produced anew phase of hydroxypyromorphite [132, 188], this mechanism is expressed as:
Ca10 (PO4)6 (OH)2 + x Pb2+ x Ca2+ + Ca10-xPbx (PO4)6(OH)2
However , this phase is not detected in the present study this may be attributed to under limit of XRD
Table 14: Ca2+ ion percentage in CA and CAHA(I) before and after metal ion uptake.
Character
Ca2+ % before uptake
Ca2+ % after uptake
Wt %
+Pb2+ (Wt %)
+ Fe3+ (Wt %)
CA 11.05 5.69 1
CAHA(I) 28.35 15.61 23.43
Also the uptake process may be occurs during chelation bonding of targeted metal ions with two carboxylic groups of alginate and one or two OH sites of the alginate ring (Fig 25) [189]. In this case metal ions may forms complexes with two adjacent alginate rings. Here,‘‘adjacent’’ means either two neighbor alginate rings of a single polymeric chain (intramolecular chelation) or two rings from two parallel chains (intermolecular chelation) [189].
Figure 25: potential active sites of CA which may bonds with targeted metal ion
a) CA + Fe+3
Element Wt % At %
C K 15.77 32.25
O K 27.57 42.32
Si K 0.41 0.35
Cl K 0.62 0.43
Ca K 1 0.61
Fe K 54.64 24.03
Total 100 100
b) CAHA(I) + Fe+3
Element Wt
% At
%
C K 7.60 12.77
O K 36.54 55.02
Al k 0.37 0.33
P K 14.03 12.91
Ca K 23.43 11.52
Fe K 18.03 7.78
Total 100 100
Figure 26: SEM images and EDX analysis of CA and CAHA(I) after iron removal
a) CA + Pb+2
Element Wt % At
%
C K 32.44 51.66
O K 36.01 43.05
Al k 0.30 0.21
Pb M 25.56 2.36
Ca K 5.69 2.71
Total 100 100
b) CAHA(I) + Pb+2
Element Wt % At %
C K 9.53 18.79
O K 32.94 48.74
Al k 0.37 0.32
P K 15.46 11.81
Pb M 8.99 1.03
Ca K 32.71 19.31
Total 100 100
Figure 27: SEM images and EDX analysis of CA and CAHA(I) after lead removal.
3.3. Metal ion uptake by adsorption process.
3.3.1. Effect of contact time on the adsorption process.
3.3.1.1 Calcium Alginate (CA):
The effect of contact time on the adsorption capacity of CA for Pb2+ ( natural pH of 5.7), Pb2+ (pH 4) and Fe3+ ( natural pH of 2.6) is represented in Figure 28. The equilibrium points of adsorption were attained within the first 60 min. for Pb2+ (82.3%) , 30 min. for Pb2+ (86.82%) and 30 min. for Fe3+ (94.15%) of contact time. the adsorption of Pb2+, Pb2+( pH 4) and Fe3+onto CA was rapid in the initial stages up to 30 min, and was almost same at high contact time. This behavior could be attributed to the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption process. The exponential phase for Pb2+( pH 4) and Fe3+ was found to be between 5 minutes and 30 minutes. For Pb2+ it was found to be between 5 minutes and 60 minutes. Further increase in contact time led to no significant adsorption of metal ions by CA probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches its maximum after 120 minutes (94.18%).
Figure 28: Effect of contact time on the adsorption of Pb2+( pH 5.7) , Pb+2(pH 4) and Fe3+ ions(pH 2.6) from aqueous solution by CA under experimental conditions of CA mass 0.02 g/10mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.1.2 CAB
The effect of contact time on the adsorption capacity of CAB for Pb2+ ( natural pH of 5.7), Pb2+(pH 4) and Fe3+ ( natural pH of 2.6) is represented in Figure 29. The equilibrium points of adsorption were attained within the first 60 minutes for Pb2+ (66.40%), 30 minutes for Pb2+ (pH 4)(68.2%) and 30 minutes for Fe3+ (89.42%) of contact time. the adsorption of Pb2+, Pb2+( pH 4) and Fe3+onto CAB was rapid in the initial stages up to 30 min, and was almost same at high contact time. This behavior may be due to the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption process. The exponential phase for Pb2+( pH 4) and Fe3+ was found to be between 5 minutes and 30 minutes. For Pb2+ it was found to be between 5 minutes and 60 minutes. The sorption efficiency by CAB was almost constant probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches its maximum after 120 minutes (89.6%).
Figure 29: Effect of contact time on the adsorption of Pb2+(pH=5.7)Pb+2(pH=4) and Fe3+(pH= 2.6)ions from aqueous solution by CAB under experimental conditions of CAB mass 0.02 g/10 mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.1.3 CAMK
The effect of contact time on the removcal of Pb2+ ( natural pH of 5.7), Pb2+ ( pH 4) and Fe3+ ( natural pH of 2.6) by CAMK is represented in Figure 30. The adsorption percent of metal ions was fast at initial stages and gradually become slower until the equilibrium is attained. The optimal contact time to attain equilibrium was experimentaly found to be about 30 min. for Pb2+ (69.26%) , Pb2+( pH 4) (77.7%) and Fe3+ (89.87%). the adsorption of Pb2+, Pb2+( pH 4) and Fe3+onto CAMK was rapid in the initial stages up to 30 minutes, and was almost same at high contact time. This behavior could be because of the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption processThe exponential phase for Pb2+, Pb2+( pH 4) and Fe3+ was found to be between 5 minutes and 30 minutes. Further increase in contact time led to no significant adsorption of metal ions by CAMK probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches its maximum after 120 minutes (90%).
Figure 30: Effect of contact time on the adsorption of Pb2+ (pH= 5.7), Pb+2( pH= 4) and Fe3+ (pH=2.6) ions from aqueous solution by CAMK under experimental conditions of CAMK mass 0.02 g/10 mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.1.4. CAHA(I)
The effect of contact time on the adsorption capacity of CAHA(I) for Pb2+ (natural pH of 5.7), Pb2+ (pH 4) and Fe3+ ( natural pH of 2.6) is represented in Figure (31). As it can be seen in fig.22, the equilibrium points of adsorption were attained within the first 60 minutes for Pb2+ (80.78%), 30 minutes for Pb2+ (pH 4) (86.62%) and 30 minutes for Fe3+ (99.33%) of contact time. the adsorption of Pb2+, Pb2+( pH 4) and Fe3+onto CAHA(I) nanopowders was rapid in the initial stages up to 30 min, and was almost same at high contact time. This behavior could be attributed to the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption process. The exponential phase for Pb2+( pH 4) and Fe3+ was found to be between 5 minutes and 30 minutes. For Pb2+ it was found to be between 5 minutes and 60 minutes. The sorption efficiency by CAHA(I) was almost constant probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches its maximum after 120 minutes (99.34%)
Figure 31: Effect of contact time on the adsorption of Pb2+ (pH= 5.7), Pb+2(pH= 4) and Fe3+ ions (pH= 2.6) from aqueous solution by CAHA(I) under experimental conditions of CAHA(I) mass 0.02 g/10 mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.1.5. CAHA (II)
The adsorption efficiency of Pb2+ ( natural pH of 5.7), Pb2+ (pH 4) and Fe3+ ( natural pH of 2.6) by using CAHA (II) was tested at different contact time. Asit can be seen in figure 32, the equilibrium points of adsorption were attained within the first 30 minutes for Pb2+ (74.08%) and Pb2+ (pH 4) (84.67%) while for Fe3+ were attained within the first 10 minutes (98.65%) of contact time. the adsorption of Pb2+, Pb2+( pH 4) and Fe3+onto CAHA(II) nanopowder was rapid in the initial stages up to 30 min, and was almost same at high contact time. This behavior could be attributed to the availability of a large number of active sites for rapid surface metal ions binding during the initial stages of adsorption process. The exponential phase for Pb2+and Pb2+( pH 4) was found to be between 5 minutes and 30 minutes. For Fe3+ it was found to be between 5 minutes and 10 minutes. Further increase in contact time led to no significant adsorption of metal ions by CAHA (II) probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches 98.78% after 60 minutes and its maximum after 120 minutes (98.81%).
Figure 32: Effect of contact time on the adsorption of Pb2+ (pH= 5.7), Pb+2(pH= 4) and Fe3+ ions (pH= 2.6) from aqueous solution by CAHA (II) under experimental conditions of CAHA (II) mass 0.02 g/10 mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.1.6. CAD
The effect of contact time on the adsorption capacity of CAD for Pb2+ ( natural pH of 5.7), Pb2+(pH 4) and Fe3+ ( natural pH of 2.6) is represented in Figure 33. The equilibrium points of adsorption were attained within the first 30 minutes for Pb2+ (68.96%), Pb2+( pH 4) (63.23%) and Fe3+ (86.72%) of contact time. The adsorption of Pb2+, Pb2+(pH 4) and Fe3+onto CAD was rapid in the initial stages up to 30 min, and was almost same at high contact time. This behavior could be attributed to the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption process. The exponential phase for Pb2+, Pb2+( pH 4) and Fe3+ was found to be between 5 minutes and 30 minutes. The sorption efficiency by CAD was almost constant probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches its maximum after 120 minutes (86.95%).
Figure 33: Effect of contact time on the adsorption of Pb2+ (pH= 5.7), Pb+2(pH= 4) and Fe3+ ions (pH= 2.6) from aqueous solution by CAD under experimental conditions of CAD mass 0.02 g/10 mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.1.7. CAG
The sorption efficiency exhibited by of CAG for Pb2+ ( natural pH of 5.7), Pb2+ (pH 4) and Fe3+ ( natural pH of 2.6) is depicted in Figure 34. The equilibrium points of adsorption were attained within the first 30 minutes for Pb2+ (83.43%), Pb2+( pH 4) (87.93%) and Fe3+ (91.11%) of contact time. The adsorption of Pb2+, Pb2+( pH 4) and Fe3+onto CAG was rapid in the initial stages up to 30 min, and was almost same at high contact time. This behavior could be due to the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption processThe exponential phase for Pb2+, Pb2+( pH 4) and Fe3+ was found to be between 5 minutes and 30 minutes.. The sorption efficiency by CAG was almost constant probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. For Fe3+ the adsorption capacity reaches 91.55% after 60 min. and reaches its maximum after 120 minutes (91.81%).
Figure 34: Effect of contact time on the adsorption of Pb2+ (pH= 5.7), Pb+2(pH 4) and Fe3+ ions (pH= 2.6) from aqueous solution by CAG under experimental conditions of CAG mass 0.02 g/10 mL, 100 mg/L metal ion concentration and temperature of 25±1oC.
3.3.2. Effect of pH
The effect of pH on the adsorption is performed only for Pb2+ because of Fe3+ solution was stable only at pH lower than 3. The study is achieved with two pHs values (4 and 5.7) the original solution of lead was at pH=5.7 and pH=4 at 5-120 minutes contact time, 20 mg dosage of the different composites and 100 ppm of metal solution at 25±1°C. As can be seen in Figures 28-34, the adsorption efficiency of Pb2+ at pH=4 is higher than that of pH=5.7 for all the composites.
3.3.3. Effect of adsorbent dosage on metal ion adsorption.
The experimental results of the adsorption of Pb2+ on CA ( as astandard model ) as afunction of adsorbent dosage 10, 15 and 20 mg/10 mL, initial Pb2+ concentration of 100 mgL-1, natural pH of 5.7, temperature 25oC at the optimal contact time (30 min) and interval contact time ( 5-30 min) are shown in fig. 36 and 37 respectively. As can be seen in Figures 35 and 36, the Pb2+ adsorption percent rapidly increased with the increase in the adsorbent dosage . this can be attributed to higher adsorbent dosage due to the increased surface area providing more adsorption sites available which gave rise to higher removal of lead [190].
Figure 35: Effect of CA dosage on Pb2+ at contact time 30 min. , initial concentration of 100 mgL-1, natural pH of 5.7 and temperature 25±1oC.
Figure 36: Effect of CA dosage on Pb2+ at contact time 5-30 min. , initial concentration of 100 mgL-1, natural pH of 5.7 and temperature 25±1oC.
3.4. kinetics studies of the adsorption process.
The kinetic study is useful to predict the adsorption rate which is very important in modeling and designing of the adsorption process [191]. The pseudo-first rate equation of lagergren and pseudo-second order kinetic model, as the most widely used models, are used to evaluate the mechanism of adsorption process.
3.4.1. The pseudo first-order model:
The linear form of the pseudo first-order kinetic rate equation is given as follows:
ln (qe –qt ) = ln qe – k1 t
where qt is the amount of metal ion adsorbed (mg g-1) at time (t), qe is the maximum adsorption capacity (at equilibrium) (mg g-1) and k1(min−1) is the rate constant for pseudo–first–order sorption, qe is the maximum adsorption capacity (at equilibrium) (mg g-1). The kinetic of adsorption are evaluated at an initial concentration of 100 mg/L for Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+(pH=2.6), adsorbent dosage of 0.02 g/10 mL and temperature of 25oC. Plot of ln (qe –qt ) vs t is drawn as shown in Figures 37-43.
The rate constant at equilibrium (qe) and regression coefficient (R2) obtained from the plots of pseudo-first rate equation of adsorbed Pb2+ and Fe3+ at equilibrium (qe) for all the adsorbents are given in Tables 15,16 and 17 respectively. As it can be seen in Figures (37-43) and Tables (15-17), The regression coefficient does not close to unity. Also, the values of qe obtained from pseudo-first order equation for all the adsorbent are different and not matched notably with the experimental qe value.
Figure 37: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CA at experimental conditions of 100 mg/L metal ion concentration, CA mass 0.02 g/10 mL and 25±1oC.
Figure 38: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAB at experimental conditions of 100 mg/L metal ion concentration, CAB mass 0.02 g/10 mL and 25±1oC.
Figure 39: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAMK at experimental conditions of 100 mg/L metal ion concentration, CAAMK mass 0.02 g/10 mL and 25±1oC.
Figure 40: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+( pH= 4) and Fe3+ (pH=2.6) adsorbed on CAHA(I) at experimental conditions of 100 mg/L metal ion concentration, CAHA(I) mass 0.02 g/10 mL and 25±1oC.
Figure 41: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAHA(II) at experimental conditions of 100 mg/L metal ion concentration, CAHA(II) mass 0.02 g/10 mL and 25±1oC.
Figure 42: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+(pH=2.6) adsorbed on CAD at experimental conditions of 100 mg/L metal ion concentration, CAD mass 0.02 g/10 mL and 25±1oC.
Figure 43: Pseudo first-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAG at experimental conditions of 100 mg/L metal ion concentration, CAG mass 0.02 g/10 mL and 25±1oC.
3.4.2. The pseudo second-order model:
The linear form of the pseudo second-order kinetic rate equation is given as follows:
t/qt = 1/k2qe2 + t/qe
where qt is the amount of metal ion adsorbed (mg g-1) at time (t), qe is the maximum adsorption capacity (at equilibrium) (mg g-1) and k2 (g mg-1 min-1) is the rate constant for the second-order sorption. from plots of t/qt versus t for the second–order reactions Figures 44-50, the k2 and qe values are calculated by using the values of intercept and slope as summarized in Tables (15, 16, 17).
The rate constant and regression coefficient (R2) obtained from the plots of pseudo-second rate equation of adsorbed Pb2+ (pH=5.7), Pb2+( pH 4) and Fe3+ (pH=2.6) at equilibrium (qe) for all the adsorbents are given in Tables 15,16 and 17 respectively . from the linear plots, the qe,experimental and the qe,calculated values are very close to each other, and also, the calculated coefficients of determination, R2, are close to unity.
Table 15: Experimental and calculated parameters of pseudo-first and second order kinetic models of Pb2+ (natural pH of 5.7) adsorbed on CA and CACs powder.
Qe experimental Pseudo first order Pseudo second order
qe calculated
K1
R2 qe calculated
K2
R2
CA
41.15
3.412
0.0769
0.9183
41.15
1.91x10-6
0.9997
CAB
33.2
5.77
0.03538
0.4826
33.69
3.4x10-5
0.9923
CAMK 34.89 11.881 0.12773 0.9636 36.16 2.41x10-5 0.9987
CAHA(I) 40.39 16.29 0.11735 0.9438 42.19 2.06x10-5 0.9984
CAHA(II)
37.09 35.55 0.19646 0.8367 39.65 3.21x10-5 0.9951
CAD 34.87 27.91 0.1423 0.8184 38.13 6.22x10-5 0.9778
CAG 42.07 11.189 0.10384 0.6573 43.3 1.33x10-5 0.9981
Table 16 . Experimental and calculated parameters of pseudo-first and second order kinetic models of Pb2+ (pH=4) adsorbed on CA and CACs powder.
qe Experimental Pseudo first order Pseudo second order
qe calculated
K1
R2 qe
calculated
K2
R2
CA 43.45 9.698 0.1718 0.9431 44.09 4.7x10-6 0.9997
CAB
34.55
13.06
0.1026
0.9001
35.95
3.63x10-5
0.9967
CAMK
38.92
27.24
0.1758
0.8219
40.7
2.34x10-5
0.9978
CAHA(I)
43.39
32.13
0.1921
0.9459
45.74
1.53x10-5
0.9955
CAHA(II) 42.45 15.99 0.1938 0.9859 44.3 1.47x10-5 0.9974
CAD
32.21
14.73
0.107
0.9513
34.16
5.62x10-5
0.9944
CAG
44.1
21.03
0.1394
0.6528
44.62
1.18x10-5
0.9932
Table 17: Experimental and calculated parameters of pseudo-first and second order kinetic models of Fe3+ (natural pH of 2.6) adsorbed on CA and CACs powder.
qe experimental Pseudo first order Pseudo second order
qe calculated
K1
R2 qe
calculated
K2
R2
CA
47.09
43.46
0.2330
0.8915
49.26
9.57x10-6
0.9995
CAB
44.75
39.40
0.2171
0.9256
47.12
1.37x10-6
0.9963
CAMK
44.95
29.695
0.2163
0.8390
46.25
6.56x10-6
0.9997
CAHA(I)
49.72
1.48
0.1037
0.8751
50
2.2x10-3
0.9999
CAHA(II)
49.39
1.41
0.0747
0.7308
50
2.2x10-3
0.9999
CAD
43.4
21.32
0.1857
0.8184
44.3
9.63x10-6
0.9974
CAG
45.75
5.47
0.0938
0.6573
46.08
3.18x10-6
0.9981
Figure 44: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4)and Fe3+ (pH=2.6) adsorbed on CA at experimental conditions of 100 mg/L metal ion concentration, CA mass 0.02 g/10 mL and 25±1oC.
Figure 45: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAB at experimental conditions of 100 mg/L metal ion concentration, CAB mass 0.02 g/10 mL and 25oC.
Figure 46: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAMK at experimental conditions of 100 mg/L metal ion concentration, CAMK mass 0.02 g/10 mL and 25±1oC.
Figure 47: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAHA(I) at experimental conditions of 100 mg/L metal ion concentration, CAHA(I) mass 0.02 g/10 mL and 25±1oC.
Figure 48: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAHA(II) at experimental conditions of 100 mg/L metal ion concentration, CAHA(II) mass 0.02 g/10 mL and 25±1oC.
Figure 49: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAD at experimental conditions of 100 mg/L metal ion concentration, CAD mass 0.02 g/10 mL and 25±1oC.
Figure 50: Pseudo second-order plots of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) adsorbed on CAG at experimental conditions of 100 mg/L metal ion concentration, CAG mass 0.02 g/10 mL and 25±1oC.
from all the obtained results illustrated in Figures 37-50 and Tables 15-17, it is obvious that the regression coefficient (R2) from pseudo-second order rate equation for all the adsorbents is higher than that of the pseudo-first order model. On the basis of the regression coefficient and calculated values of adsorption capacity, the adsorption process is found to obey and exhibited best fit to the pseudo-second-order kinetic model which mean that the rate-limiting step might be chemical adsorption or chemisorption involving valency forces through exchange of electrons between the sorbate and the sorbent, also only one ion of the metal is sorbed onto two sorption sites on the sorbent surface [192, 193].
3.4.3. Prediction of adsorption rate-limiting step
There are essentially three consecutive mass transport steps associated with the adsorption of solute from the solution by an adsorbent. These are (1) film diffusion, (2) intraparticle or pore diffusion, and (3) sorption into interior sites. The third step is very rapid and hence, film and pore transports are the major steps controlling the rate of adsorption [179, 194].
The most commonly used technique for identifying the mechanism involved in the adsorption process is by fitting an intraparticle diffusion plot [195]. The amount of metal ions adsorbed (qt) at time (t), is plotted against the square root of t (t1/2), according to Eq. proposed by Weber and Morris as follows:
Qt = Kid t0.5 + C
where C is constant and kid is the intraparticle diffusion rate constant (mg/g min1/2), qt is the amount adsorbed at a time (mg/g), t is the time (min), and kid (mg/g min1/2) is the rate constant of intraparticle diffusion. Due to the varying extent of adsorption in the initial and final stages of the experiment two straight lines with different slopes are obtained (Figures 51-57).
The two regions in the qt vs. t0.5 plot suggest that the sorption process proceeds by surface sorption and intraparticle diffusion. The initial rapid uptake can be attributed to the boundary layer effects (film diffusion). After the external surface loading was completed, the intraparticle diffusion or pore diffusion takes place, The second linear part of the plot presented in Fig.51-57, corresponds to the transportation of Pb2+ and Fe3+ within CA and CACs particles [196]. The slope of the second linear portion of the plot has been defined to yield the intraparticle diffusion parameter of ki1, Ri12 (first stage) and ki2, Ri22 (second stage) are listed in Tables (18-20). On the other hand, the intercept of the plot give an idea about the thikness of boundary layer effect [197]. The larger the intercept, the greater the contribution of the surface sorption in the rate-controlling step [195].
As it can be seen in Figures (51-57), the plot indicated that the intraparticle diffusion was not the rate-controlling step because it did not pass through the origin [196]. The deviation of the straight lines from the origin may be due to the difference in the rate of mass transfer in the initial and final stages of adsorption [198]. Further, the first straight portion is attributed to a macropores diffusion process and the second linear portion can be ascribed to a micropore diffusion process [192, 199]. In addition, it is clear from Fig. (51-57) that the first stage is faster than the second one. This behaviour may be correlated with the very slow diffusion of the adsorbate from the surface film into the micropores, which are the least accessible sites for adsorption [197].
Figure 51: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+(pH=2.6) on CA.
Figure 52: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) on CAB composite.
Figure 53: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) on CAMK composite.
Figure 54: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+( pH 4) and Fe3+ (pH=2.6) on CAHA(I) composite.
Figure 55: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) on CAHA(II) composite.
Figure 56: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) on CAD composite.
Figure 57: Intraparticle diffusion plot for adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+ (pH=2.6) on CAG composite.
Table 18: The parameters of intraparticle diffusion model for adsorption of Pb2+ (pH=5.7) onto CA and CACs.
Intrapartical diffusion modle
Ki1
(mg/g m0.5)
Ki2
(mg/g m0.5)
Ri12
Ri22
Intercept i1
Intercept i2
CA
0.617 0.071 0.9541 0.7161 37.55 40.50
CAB
1.460 0.13 0.8094 0.7094 24.3 32.07
CAMK
1.682 0.049
0.9365 0.6601 25.94 34.31
CAHA(I)
3.53 0.159 0.9989 0.6689 23.43 38.85
CAHA(II)
2.087 0.019 0.9917 0.8957 25.62 36.93
CAD
3.77 0.15 0.7123 59.22 15.48 33.45
CAG
1.36 0.12 0.9615 0.8696 33.92 41.05
Table 19: The parameters of intraparticle diffusion model for adsorption of Pb2+ (pH=4) onto CA and CACs.
Intrapartical diffusion modle
Ki1
(mg/g m0.5)
Ki2
(mg/g m0.5)
Ri12
Ri22
Intercept i1
Intercept i2
CA
1.11 0.066 0.9056 0.4133 38.07 42.87
CAB
1.71 0.10 0.9716 0.7251 33.6 24.54
CAMK
1.88 0.021 0.9745 0.9251 28.48 38.74
CAHA(I)
3.21 0.086 0.8603 0.3367 27.94 42.06
CAHA(II)
2.01 0.039 0.9556 0.9346 31.56 42.12
CAD
2.25 0.14 0.9119 0.7920 19.74 30.89
CAG
1.43 0.025 0.7045 0.4308 34.97 43.84
Table 20: The parameters of intraparticle diffusion model for adsorption of Fe3+ (pH=2.6) onto CA and CACs.
Intrapartical diffusion modle
Ki1
(mg/g m0.5)
Ki2
(mg/g m0.5)
Ri12
Ri22
Intercept i1
Intercept i2
CA
2.75 0.097 0.9892 0.108 33.98 46.16
CAB
3.79 0.11 0.7902 0.1968 27.28 43.7
CAMK
1.63 0.012 0.9593 0.9589 36.36 44.85
CAHA(I)
0.3 0.0017 0.6592 0.3204 48.14 49.65
CAHA(II)
0.17 0.013 0.9673 0.5278 48.37 49.25
CAD
1.33 0.020 0.9937 0.9134 35.9 43.24
CAG
0.62 0.062 0.9351 0.7898 41.65 45.22
from Tables (18-20) it can be conclouded that for adsorption of Pb2+(pH=5.7), Pb2+ (pH=4) and Fe3+ (pH=2.6) onto CA and CACs, the calculated values of ki1 were higher than that of ki2. The reason could be as circumscription of the available vacant space for diffusion in them, so of pore blockage. The values of the correlation regression coefficients characterizing the applicability of the intraparticle diffusion model (Ri12, Ri22) were lower than that of R2 (pseudo second order), but commensurable with R12 ( pseudo first order). Actually, the three models stated above could describe the proposed sorption process to a definite extend, but they could not predict the high rate of adsorption during the first minutes of the process. Probably, the initial stages are controlled by external mass transfer or surface diffusion, followed by chemical reaction or a constant-rate stage, and diffusion causing gradual decrease of the process rate [192]
3.5. Adsorption isotherms:
Adsorption isotherm studies are necessary for illustrating the adsorption process at equilibrium conditions. An adsorption isotherm is characterized by certain constants which express the adsorbent affinity and can also be used for finding the adsorption capacity of the sorbent. The adsorption of iron and lead from polluted water using CA and CACs could be assumed to have abehavior fitting with the isothermal adsorption model in which the adsorbate keeps a dynamic equilibrium between the adsorotion and desorption at afixed temperature [178, 200].
Two most widely used mathematical models Langmuir and Freundlich adsorption isotherms are adopted for expressing the quantitative relationship between the extent of sorption and the residual solute concentration. Langmuir adsorption isotherm assumes monolayer coverage of adsorabate over ahomogeneous adsorbent surface and the adsorption of each molecule onto the surface has the same activation energy of adsorption.
Freundlich adsorption isotherms assumes aheterogeneous surface with anon-uniform distribution of heat of adsorption over the surface with the possibility of the multilayer adsorption [178, 179]. The maximum metal ions adsorption capacities are determined by analyzing the experimental data for heavy metal adsorption onto CA and CAHA(I) [201], as they provide the higher removal effeciency. The data of Pb2+ and Fe3+ adsorption by CA and CAHA(I) are examined in accordance with langmiur adsorption isotherm models whose linearized equation was:
Ce/qe =Ce/qm + 1/ KL qm
where qm (mg /g) and KL (L /mg) are constants in Langmuir’s equation which are referred to the maximum adsorption capacity corresponding to complete monolayer coverage and the Langmuir model constant that is indirectly related to the energy of adsorption. Also qe and Ce parameters represent the equilibrium adsorption capacity and the equilibrium concentration heavy metal ion that is remaining in solution, respectively. qe is calculated as follows:
qe =((Co-Ce)V)/m
where Co is the initial metal ion concentration (mg/L), Ct is the equilibrium concentration of adsorbate (mg/L) (mg/L), V is the initial solution volume (L) and m is the adsorbate dose (g). A plot of Ce/qe vs. Ce (Fig. 58-61) gives a linear trace with a slope of 1/qm and intercept of 1/ KL qm. A further analysis of the Langmuir equation can be made on the basis of a dimensionless equilibrium parameter, RL, also known as the separation factor,
given by
RL=1\(1+ KLCe)
where Ce is equilibrium liquid phase concentration of the solute at which adsorption is carried out. The value of RL lies between 0 and 1 for favorable adsorption, while RL > 1 represents unfavorable adsorption, and RL = 1 represents linear adsorption, while the adsorption process is irreversible if RL = 0 [179, 202].
Also the obtained data are examined in accordance with the Freundlich equation which is expressed as
qe =log KFCe1/n
where KF is the measure of adsorption capacity and n is the adsorption intensity linear form of Freundlich
log qe =log KF + (1/n) log Ce
where qe is the amount adsorbed (mg/g), Ce is the equilibrium concentration of adsorbate (mg/L), and KF and n are the Freundlich constants related to the adsorption capacity and adsorption intensity, respectively. A plot of log qe vs. log Ce (Fig. 58-61) gives a linear trace with a slope of 1/n and intercept of log KF. The 1/n value in the range of 0 and 1 is a predicting of adsorption intensity of metal ion onto the adsorbent and the type of isotherm to be irreversible (1/n=0), favourable (0<1/n<1) and unfavourable (1/n >1) [197], the 1/n value determine the surface heterogeneity, becoming more heterogeneous as its value gets closer to zero. In addition, the value of n varies with the heterogeneity of the adsorbent, if n < 10 and n > 1 indicating the adsorption process is favorable [193, 203].
The isotherm parameters and correlation coefficients calculated for the adsorption of Pb2+ and Fe3+ using CA and CAHA(I) are listed in Tables (21,22).
Table 21: Isotherm parameters and correlation coefficients calculated for the adsorption of Pb2+ using CA and CAHA(I).
Adsorbent
Langmuir Isotherm
Freundlich Isotherm
qmax(mg/g) KL R2 KF 1/n N R2
CA
51.78
1.264
0.9353
18.62
0.377
2.64
0.6028
CAHA(I)
52.99
0.702
0.8652
16.89
0.3450
2.89
0.9555
Table 22: Isotherm parameters and correlation coefficients calculated for the adsorption of Fe3+ using CA and CAHA(I).
Adsorbent
Langmuir Isotherm
Freundlich Isotherm
qmax(mg/g) KL R2 KF 1/n N R2
CA 66.53 0.0816 0.9361 15.84 0.2602 3.84 0.9656
CAHA(I) 113.63 0.1560 0.7675 24.35 0.3334 2.99 0.9178
from Table 21, it can be concluded that for adsorption of lead on CA, the Langmuir isotherm (R2 > 0.93) fitted the experimental results better than those of the Freundlich isotherm (R2 > 0.60) as reflected with the correlation coefficient, indicating the homogenous feature presented on the CA surface and demonstrates the formation of monolayer coverage of the lead ions on the CA surface, the adsorption is localized, all active sites of surface have similar energies and no interaction between adsorped molecules. Moreover, the value of RL was 0.0003. This also suggests an irreversible adsorption between CA and Pb2+ ions [202].
On the other hand, for CAHA, it can be stated that the Freundlich isotherm (R2 > 0.95) fitted the experimental results comparable to the Langmuir isotherm (R2 > 0.86), indicating that the adsorbed amount increased with initial concentration. The slope 1/n provides information about surface heterogeneity and surface affinity for the solute. As a higher value of 1/n (0.34) is obtained, it corresponds to the greater heterogeneity of the adsorbent surface. Furthermore, the value of n > 1 obtained from the Freundlich isotherm indicating (2.8), that this process is also favorable [203] and heterogeneous sorption. The maximum adsorption capacities of the Pb2+ ions are found to be 51.78 and 52.99 mg/g for CA and CAHA(I), respectively.
from Table 22, it can be stated that for adsorption of iron on CA and CAHA(I), the Freundlich isotherm (R2 > 0.96), (R2 > 0.91) fitted the experimental results comparable to the Langmuir isotherm (R2 > 0.93), (R2 > 0.76) for CA and CAHA(I) respectively. The slope 1/n provides information about surface heterogeneity and surface affinity for the solute. As a higher value of 1/n is obtained, it corresponds to the greater heterogeneity of the adsorbent surface. Furthermore, the value of n > 1 obtained from the Freundlich isotherm (3.8, 2.9), indicating that this process is also favorable and heterogeneous sorption.
Firure 58: Langmuir and Freundlish isotherms for adsorption of Pb2+ on CA Powder.
Firure 59: Langmuir and Freundlish isotherms for adsorption of Pb2+ on CAHA(I) Powder.
Firure 60: Langmuir and Freundlish isotherms for adsorption of Fe3+ on CA Powder.
.
Firure 61: Langmuir and Freundlish isotherms for adsorption of Fe3+ on CAHA(I) Powder
SUMMARY AND CONCLUSION
The present thesis comprises of there chapters:
Chapter (1) includes introduction and literature review that focused on the heavy metals found in industrial effluents as hazardous pollutants that may affect hardly the environment.
This chapter also includes various technologies which have been used to remove metal ions. Especial attention is given to two types of heavy metals ions (Pb2+ and Fe3+), their abundance, use in several industries and harmful effect on the environment especially to aquatic systems.
The literature review also focused on calcium alginate and its biomedical applications, also focused on bentonite, metakaolin, diatomeous silica, gelatin and hydroxy apatite (produced from biowastes) which form composites with calcium alginate and their role in heavy metals removal.
Chapter (2) includes experimental and methods which focused on types of chemical used, methods of preparation of calcium alginate, hydroxy apatite and calcium alginate composites and characterization tecniques such as X-Ray Diffraction (XRD), Fourier Transformer Infrared (FTIR), Scanning Electron Microscope(SEM) and Energy Dispersive Analysis X-ray (EDX). Preparation of heavy metal solutions with the assessment of their concentrations before and after adsorption process using Perkin Elmer 2380 Atomic Absorption Spectro Photometer. The six synthesized composite beside calcium alginate used in the present study are listed in the following table:
Composite Compound Source of raw material Abbrev.
1 Calcium alginate Oxford Lab. Reagent CA
2 Calcium alginate-Bentonite Bentonite (Abu Zaabal Fertilizer & Chemicals Co. CAB
3 Calcium alginate-Metakaolin Metakaolin (Sinai Peninsula) CAMK
4 Calcium alginate-HAP(1) HAP (egg shell calcined at 900oC) CAHA(I)
5 Calcium alginate-HAP(2) HAP (bovin bone calcined at 1000oC) CAHA(II)
6 Calcium alginate-Diatomeous Diatomeous (Kazakhstan) CAD
7 Calcium alginate-Gelatin Oxford Lab. Reagent CAG
Chapter (3) is concerned with the results and discussion that includes characterization of the prepared composites before and after heavy metal uptake using XRD, FTIR, SEM and EDX techniques. It also include the adsorption process of heavy metal uptake and kinetics studies of the adsorption process.
I. Fourier Transform Infrared (FTIR) analysis:
The FTIR spectra of calcium alginate showed important absorption bands regarding hydroxyl, ether and carboxylic functional groups. Stretching vibrations of O–H bonds of alginate appeared at 3444 cm-1. In the spectra of CAG, the absorption band at around 3442 cm-1 concerned with OH stretching vibration for CA slightly broadened and shifted to a lower wave number with the blending with gelatin, suggesting the formation of an intermolecular hydrogen bond.
CA raw material, reveals asimilar spectra with CAB. A strong and sharp band is detected at 1022 cm-1 which is related to Si–OH stretching vibrations. 1034 cm-1 assigned to Si-O stretching, and 875 cm-1 is due to OH bending of the Al-Al-OH group. A similar OH bending vibration is observed for Al-Mg-OH at 842 cm-1, 690 cm-1 assigned to quartz. Also, there is a shoulder peak at 520 cm-1 (Al-O-Si bending), and 464 cm-1 (Si-O-Si bending). The band of Si-O-Si bending in CAD is the strongest and sharpest compared to CAMK and CAB.
The bands corresponding to the samples of CAHA(I) and CAHA(II) which appears at 3570 and 630-633 cm-1 (due to structural OH of hydroxy apatite) confirm the hydroxyapatite. The most intensive bands in the range of 1044 –1090 cm-1 corresponded to the triply degenerated asymmetric stretching vibrations of P-O. Otherwise the peak at 962.97 cm-1 indicates the non-degenerated asymmetric mode of PO43-. The very strong and sharp bands observed at 569-572 and 602-603 cm-1 attributed to triply degenerated bending mode of the O-P-O in PO43- group. The larges parting distance of these bands revels the crystalline phase.
II. X-ray diffraction for crystal phase detection:
The X-ray diffraction pattern of CA and CACs in the range of 2θ = 5-60o diffraction degree shows two typical peaks in 2θ =16° and 22° corresponding to calcium alginate. The XRD of CAG shows typical peaks around 12° and 21°. The samples of CAB, CAMK and CAD revealed very similar XRD pattern peaks of quartz phase.
The sample CAHA(I) and CAHA(II) revealed very similar XRD pattern peaks of 100 % HAP (reference code :01-086-1194) in the rang of 2θ =31o and 32o and d (Ao) value spacing equal 2.81 and 2.78.
XRD analysis of CAHA(II) also indicated the absence of secondary phases, such as tri calcium phosphate (TCP) or calcium oxide (CaO). In the case of CAHA(I), its diffraction pattern revealed additional phase of β-tricalcium phosphate, beside the HAP as a main phase. Hydoxyapatite prepared from bovine bone contain certain amount of carbonate (CO32-) in its lattice structure. The carbonate ions affect on the the degree of crystallinite and hence increase the bioactivity of pure HAP.
III. Scanning Electron Microscope (SEM) and Energy Dispersive X-ray (EDX)
According to the results obtained by EDX analysis, carbon, oxygen and calcium are the major constituents of all the prepared composites related to calcium alginate polymer present. EDX analyses of CA and CAG revealed that the major elements present are carbon and oxygen corresponding to polymer composition. The SEM of CAG showed a smooth and homogeneous morphology, suggesting high miscibility and blend homogeneity between calcium alginate and gelatin.
In the case of CAB, CAMK and CAD, the presence of an obvious peak related to the Si compounds is evident. According to the results obtained by EDX analysis, weigh percent of elements present indicating a large portion of the composites is composed of Si compounds which is suitable for an efficient sequestering metal cations from aqueous solution.
EDX analyses of the prepared HAP revealed that inorganic phases of bovin bone and egg shells were mainly composed of calcium and phosphorus as the major constituents with some minor components such as C, O, Na, Mg and Si. The weight and atomic percentage shows that the Ca/P ratio around 1.7 and 1.8 which is below 2 and acceptable where the ideal Ca/P ratio of HA is 1.67, in HAP composites (CAHA(I) and CAHA(II)) this ratio increased than 2 due to the excess amount of Ca crosslinkage of calcim alginate presents.
IV. FTIR Analysis after metal ion uptake:
FTIR spectra of CA and CAHA(I) after targeted metal ions uptake ( Pb2+, Fe3+ ) showed that there is no significant change in peak positions after metal ion uptake and also no new absorption peaks are detected . There are alittle peak shifts which may be attributed to the corporation and substitution of metal ion in the lattice structure of CA and CAHA(I).
V. X-ray diffraction for crystal phase detection after metal ion uptake:
XRD patterns of CA and CAHA(I) after Pb2+ and Fe3+ metal ions removal did not revealed any new phases. supported the proposal that Pb2+ and Fe3+ ions uptake was not dependent on dissolution/precipitation mechanisms. Pb2+ and Fe3+ ions removal may be occurs by adsorption mechanisms like surface complexation or ionic exchange. The XRD patterns showed some changes in their relative intensities and crystal sizes. Also, ther are some little shifts in d- spacing values, this may be due to the ion exchange between Ca2+ and metal cations of Pb2+ and Fe3+ in lattice structure of CA and CAHA(I) Powder.
VI. Scanning Electron Microscope (SEM) and Energy Dispersive X-ray (EDX) after metal ion uptake:
SEM of CA and CAHA(I) after metal ion uptake revealed some changes in morphology and microstructure. Morever EDX results indicated the presence of Pb2+ and Fe3+ ions with CA and CAHA(I). the peak of Fe3+ is more intense than that of Pb2+ and this confirms, the higher value of Fe3+ ions uptake by CA and CAHA(I) comparing with Pb2+ ions.
The weight and atomic percentage of Ca2+ ions in CA and CAHA(I) after metal ion uptake was less than its value before metal ion removal. The decrease in calcium percentage may be attributed to ion exchange between targeted metal ions and; (1)calcium ions of calcium alginate in CA and CAHA(I), (2) or Ca ions of HA present in CAHA(I). This ion exchange mechanism between Pb2+ ions (as example) and Ca2+ ions of HA produced anew phase of hydroxypyromorphite. However , this phase is not detected in the present study this may be attributed to under limit of XRD.
Also the uptake process may be occurs during chelation bonding of targeted metal ions with two carboxylic groups of alginate and one or two OH sites of the alginate. In this case metal ions may forms complexes with two adjacent alginate rings. Here,‘‘adjacent’’ means either two neighbor alginate rings of a single polymeric chain (intramolecular chelation) or two rings from two parallel chains (intermolecular chelation).
VII. Metal ion uptake by adsorption process.
1- Effect of contact time on the adsorption process.
The effect of contact time on the adsorption capacity of CA and CACs for Pb2+ ( natural pH of 5.7), Pb2+ (pH=4) and Fe3+ ( natural pH of 2.6) indicated that, the equilibrium points of adsorption are attained within the first 10 – 60 min. of contact time with different removal efficiency and slightly similar exponential phase as listed in table below. the adsorption of Pb2+(pH=5.7), Pb2+(pH=4) and Fe3+(pH=2.6)onto CA and CACs powder was rapid in the initial stages and was almost same at high contact time. This behavior could be attributed to the availability of a large number of active sites for rapid surface metal ions binding during the first stage of adsorption process. Further increase in contact time led to no significant adsorption of metal ions by the adsorbents probably due to the decrease in the diffusion rate since the sites of the adsorbent are covered with metal ions. As it can be seen in the following table, CAG is the higher efficiency for lead (Pb2+) removal then CAHA(I) and CA. On the other hand, CAHA(I) is the higher efficiency for iron (Fe3+) removal then CAHA(II) and CA.
Equilibrium point time (min.) Removal effeciency(%) Exponential phase time (min.)
Pb2+
(pH=5.7)
Pb2+
(pH=4)
Fe3+
(pH=2.6)
Pb2+
(pH=5.7)
Pb2+
(pH =4)
Fe3+
(pH=2.6)
Pb2+
(pH=5.7)
Pb2+
(pH=4)
Fe3+
(pH=2.6)
CA
60
30
30
82.3
86.82
94.15
5-60
5-30
5-30
CAB
60
30
30
66.4
68.2
89.42
5-60
5-30
5-30
CAMK
30
30
30
69.26
77.7
89.87
5-30
5-30
5-30
CAHA(I)
60
30
30
80.78
86.62
99.33
5-60
5-30
5-30
CAHA(II)
30
30
10
74.08
84.67
98.65
5-3
5-30
5-10
CAD
30
30
30
68.96
63.23
86.72
5-30
5-30
5-30
CAG
30
30
30
83.43
87.93
91.11
5-30
5-30
5-30
2- Effect of pH
The effect of pH on the adsorption is performed only for Pb2+ because of Fe3+ solution was stable only at pH lower than 3. The study is achieved with two pHs values (4 and 5.7) the original solution of lead is at pH=5.7 and pH=4 at 5-120 minutes contact time, 20 mg dosage of the different composites and 100 ppm of metal solution at 25±1°C. The adsorption efficiency of Pb2+ at pH=4 is higher than that of pH 5.7 for all the composites.
3- Effect of adsorbent dosage on metal ion adsorption.
The experimental results of the adsorption of Pb2+ on CA ( as astandard model ) as afunction of adsorbent dosage 10, 15 and 20 mg/10 mL, initial Pb2+ concentration of 100 mgL-1, natural pH of 5.7, temperature 25oC at the optimal contact time (30 min) and interval contact time ( 5-30 min) showed that, the Pb2+ adsorption percent rapidly increased with the increase in the adsorbent dosage . this can be attributed to higher adsorbent dosage due to the increased surface area providing more adsorption sites available which gave rise to higher removal of lead.
VIII. kinetics studies of the adsorption process.
The kinetic study is useful to predict the adsorption rate which is very important in modeling and designing of the adsorption process. The kinetic of adsorption are evaluated at an initial concentration of 100 mg/L for Pb2+(pH=5.7),Pb2+( pH 4)and Fe3+(pH=2.6), adsorbent dosage of 0.02 g/10 mL and temperature of 25oC. By appling the pseudo-first rate equation of lagergren, it is clear that the regression coefficient does not close to unity. Also, the values of qe obtained from pseudo-first order equation for all the adsorbent are different and not matched notably with the experimental qe value. from the linear plots of pseudo-second rate equation of lagergren, the qe,experimental and the qe,calculated values are very close to each other, and also, the calculated coefficients of determination, R2, are close to unity
from all the obtained results, it is obvious that the regression coefficient (R2) from pseudo-second order rate equation for all the adsorbents was higher than that of the pseudo-first order model. On the basis of the regression coefficient and calculated values of adsorption capacity, the adsorption process was found to obey and exhibited best fit to the pseudo-second-order kinetic model which is mean that the rate-limiting step might be chemical adsorption or chemisorption involving valency forces through exchange of electrons between the sorbate and the sorbent, also only one ion of the metal is sorbed onto two sorption sites on the sorbent surface.
IX. Prediction of adsorption rate-limiting step
There are essentially three consecutive mass transport steps associated with the adsorption of solute from the solution by an ads0rbent. These are (1) film diffusi0n, (2) intraparticle or p0re diffusion, and (3) sorption into interior sites. The third step is very rapid and hence, film and pore transports are the major steps controlling the rate of adsorption. The most commonly used technique for identifying the mechanism involved in the adsorption process is by fitting an intraparticle diffusion plot proposed by Weber and Morris.
The results stated that the sorption process proceeds by surface sorption and intraparticle diffusion. The initial rapid uptake can be attributed to the boundary layer effects (film diffusion). After the external surface loading was completed, the intraparticle diffusion or pore diffusion takes place. However, the plot indicated that the intraparticle diffusion was not the rate-controlling step because it did not pass through the origin.
X. Adsorption isotherms:
Adsorption isotherm studies are necessary for illustrating the adsorption process at equilibrium conditions. Two most widely used mathematical models Langmuir and Freundlich adsorption. Langmuir adsorption isotherm assumes monolayer coverage of adsorabate over ahomogeneous adsorbent surface and the adsorption of each molecule onto the surface has the same activation energy of adsorption. Freundlich adsorption isotherms assumes aheterogeneous surface with anon-uniform distribution of heat of adsorption over the surface with the possibility of the multilayer adsorption
The results of single metal ion adsorption of Pb2+ onto CA at 25 oC can be represented well by langmiur than Freundlish model with good correlation coefficient (R2). This means that the adsorbates containing Pb2+ was adsorbed in such amanner that only one atomic layer of adsorbate can be adsorbed and distributed uniformly on the surface of the adsorbents (CA) and the adsorption of each molecule onto the surface has the same activation energy of adsorption. the value of RL was 0.0003. This also suggests an irreversible adsorption between CA and Pb2+ ions.
For iron and lead adsorbed onto CAHA and iron adsorbed into CA, it can be stated that the Freundlich isotherm well fitted the experimental results comparable to the Langmuir isotherm indicating that the adsorbed amount increased with initial concentration. The slope 1/n provides information about surface heterogeneity and surface affinity for the solute. As a higher value of 1/n is obtained, it corresponds to the greater heterogeneity of the adsorbent surface. Furthermore, the value of 1 < 1/n > 0 and the value of n > 1 obtained from the Freundlich isotherm indicating, that this process is also favorable and heterogeneous sorption.
from all the obtained results and analysis we can stated that, the uptake of Pb2+ and Fe3+ may be occurs by adsorption mechanisms like surface complexation during chelation bonding of targeted metal ions with two carboxylic groups of alginate and one or two OH sites of the alginate ring forming complexes with two adjacent alginate rings. Here,‘‘adjacent’’ means either two neighbor alginate rings of a single polymeric chain (intramolecular chelation) or two rings from two parallel chains (intermolecular chelation). or ion exchange between targeted metal ions and; (1) Calcium ions of calcium alginate in CA and CACs, (2) or Ca ions of HA present in CAHA(I) and CAHA (II).
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الملخص العربي
الرساله تنقسم الي ثلاثة فصول رئيسية :
الفصل الاول : ويتضمن المقدمة والمراجع التاريخية التي تخص العناصر الثقيلة كمُلوثات بيئية والتي تنتج بكمية كبيرة من المخلفات الصناعية والتي تؤثر بصورة خطيرة علي البيئة وصحة الانسان. وشمل هذا الفصل أيضا مختلف التقنيات المستخدمه لأزالة هذه العناصر. وتم القاء الضوء علي عنصرين من هذه العناصر وهما الرصاص والحديد من حيث مصادر تواجدها من المخلفات الصناعية وتأثيرها الضارعلي البيئة المائية. ويتضمن هذا الفصل أيضا شرح وافي لألجينات الكالسيوم وتطبيقاتها في ازالة العناصر الثقيلة, والجيلاتين وبعض الطفلات مثل البنتونيت والميتاكاولين وسيليكا الدايتموس والهيدروكسي اباتيت المحضرة من المخلفات الحيوية (مثل قشر البيض وعظام البقر) والتي تكون مركبات مع الجينات الكالسيوم للتخلص من العناصر الثقيلة.
الفصل الثاني : ويشمل الكيماويات المستخدمة , وطرق تحضير ألجينات الكالسيوم والهيدروكسي اباتيت ومركبات ألجينات الكالسيوم المختلفة وطرق التعرف عليها باستخدام قياسات حيود الاشعة السينية , الأشعة تحت الحمراء, الميكرسكوب الألكتروني الماسح للضوء وطاقة الاشعة السينية المشتتة. وكذلك طرق تحضير محاليل ايونات العناصر الثقيلة وتقديرتركيزها قبل وبعد عملية الإمتزاز باستخدام جهاز الامتصاص الذري. ويتضمن الجدول التالي مصادر المخلفات الحيوية والمواد الخام وألجينات الكالسيوم ومركباته السته المحضرة :
Composite Compound Source of raw material Abbrev.
1 Calcium alginate Oxford Lab. Reagent CA
2 Calcium alginate-Bentonite Bentonite (Abu Zaabal Fertilizer & Chemicals Co. CAB
3 Calcium alginate-Metakaolin Metakaolin (kaolin (Sinai Peninsula) calcined at 800oC CAMK
4 Calcium alginate-HAP(1) HAP (egg shell calcined at 900oC) CAHA(I)
5 Calcium alginate-HAP(2) HAP (bovin bone calcined at 1000oC) CAHA(II)
6 Calcium alginate-Diatomeous Diatomeous (Kazakhstan) CAD
7 Calcium alginate-Gelatin Oxford Lab. Reagent CAG
الفصل الثالث : ويتناول النتائج التي تم الحصول عليها ومناقشتها من حيث توصيفها قبل وبعد عملية الامتزاز باستخدام اجهزة SEM وEDX مع XRD و FTIR ويشمل ايضا دراسة عملية الامتزاز و حركية عملية الامتزاز.
التعرف علي المجموعات الوظيفية من قياسات الاشعة تحت الحمراء(i)
اظهرت تحاليل عينة الجينات الكالسيوم تواجد حزم امتصاص هامة تعود الي المجموعات الوظيفيه الخاصة ب الهيدروكسيل والكربوكسيليك والايثر. حيث تظهر مجموعة الهيدروكسيل الخاصه ألجينات الكالسيوم CAعند طول موجي حوالي 3444 cm-1 , ويظهر طيف مركب الجينات الكالسيوم مع الجيلاتين CAG أن ذروة الامتصاص عند حواليcm-1 3442 والخاصة باهتزاز مجموعة الهيدروكسيل OH لـ ألجينات الكالسيوم قد اتسعت قليلاً وتحركت إلى طول موجي أقل بالمزج مع الجيلاتين ، مما يشير إلى تكوين رابطة هيدروجينية بين جزيئية.
يوجد تشابه كبير بين طيف ألجينات الكالسيوم وطيف ألجينات الكالسيوم مع البنتونيت CAB والميتاكاولين CAMK وسيليكا الدايتاموس CAD , ففي CAB تظهر ذروة امتصاص حادة وقوية عندcm-1 1022 والخاصة باهتزاز مجموعة Si-OHوعند cm-1 1034 لمجموعة Si-O وعند 875 cm-1 بسبب وجود OH bending لمجموعةAl-Al-OH كذلك تهتز مجموعة مشابهه منOH bending , خاصة ب Al-Mg-OH تظهر عند842 cm-1 و cm-1 690 تعود الي تواجد الكوارتز. ايضا يوجد حزمة امتصاص شولدر عند 520 cm-1 تعود الي (Al-O-Si bending ) وعند 464 cm-1 خاصة ب (Si-O-Si bending). تلاحظ ايضا ان حزمة الامتصاص الخاصة ب Si-O-Si bending في عينة مركب ألجينات الكالسيوم مع الداياتاموس تكون اقوي وأحد مقارنة بعينات CAB و CAMK.
بالنسبة لعينات ألجينات الكالسيوم مع الهيدروكسي أباتيت المحضرة من قشور البيض وعظام البقر CAHA(I) و CAHA(II) فتظهر حزم امتصاص عند 3570 cm-1 و 630-633 cm-1 بسبب اهتزازمجموعة OH للهيدروكسي أباتيت. كما أظهرت النتائج ايضا أن ذروة الامتصاص الاكثر شده في المدي من 1044 cm-1 الي 1090 cm-1 ترجع الي اهتزازات الرابطة P-O في مجموعة الفوسفات وتكون تللك الاهتزازات متماثلة , اما عند 962.97 cm-1 فتكون تلك الاهتزازات لمجموعة الفوسفات غير متماثلة. حزم الامتصاص القوية جدا والحادة والتي تظهر عند 569-572 cm-1 و عند 602-603 cm-1 فتعزي الي اهتزازات مجموعات O-P-O في مجموعة الفوسفات PO43-.
(ii) قياسات حيود الاشعه السينية :
بينت التحاليل باستخدام X-ray diffraction لعينات ألجينات الكالسيوم ومركباته في في المدي 5-60o=θ2 تواجد عدد ذروتين خاصة بالجينات الكالسيوم في نطاق =16 o , 22 oθ2 واظهرت النتائج ايضا تواجد الجيلاتين في عينة CAG في نطاق =12 o , 21 oθ2 وتظهر نتائج تحاليل عينات CAB, CAD, CAMK تشابه كبيربينهم و تكون الطور الخاص بالكوارتز.
كما تظهر النتائج تشابه كبير لعينات الجينات الكالسيوم مع الهيدروكسي اباتيت المحضر من قشور البيض CAHA(I) و CAHA(II) ذات الهيدروكسي اباتيت المحضر من عظام الابقار وتكون طور الهيدروكسي اباتيت طبقا للكود المرجعي 01-086-1194 في نطاق =31 o , 32 oθ2 عند قيم مسافات تباعد (d-spacing) مساوية 2.81 و 2.78انجستروم , ولا تحتوى عينات CAHA(II) على أي أطوار اخرى مثل أكسيد الكالسيوم CaO أوفوسفات الكالسيوم في حين ان عينات CAHA(I) تظهر تواجد طور فوسفات الكالسيوم بجانب طور الهيدروكسي اباتيت كطور رئيسي وتظهر نتائج تحاليل عينة الجينات الكالسيوم مع الهيدروكسي اباتيت المحضرة من عظام الابقار CAHA(II) انها تحتوي علي كمية من الكربونات في الشبكة البلورية. وتوفر ايونات الكربونات زيادة في النشاطية الحيوية للهيدروكسي أباتيت كما تؤثر في درجة التبلور للمركب.
(iii) التعرف على الشكل المورفولوجى والتحليل النوعي والكمى للعينات المحضرة بأستخدام قياسات الميكروسكوب الإلكترونى الماسح للضوء وطاقة الأشعة السينية المشتتة.
أظهرت صور الميكروسكوب الإلكتروني لجميع العينات تكتل بلوري للجزيئات مع شكل غير منتظم
نسبيا ، مع أحجام مختلفة من البللورات وتوضح الصور الشكل المورفولوجي المتجانس والناعم لعينة الجينات الكالسيوم مع الجيلاتين مشيرة الي الخلط المتجانس بينهما. وأثبتت خرائط EDX لجميع العينات تواجد اشعاع Kα لعناصر الكربون والاكسجين والكالسيوم والمفترض تواجدها في بوليمرات ألجينات الكالسيوم وتظهر النتائج ايضا في العينات المحتوية علي الهيدروكسي أباتيت ان الأطوار غير العضوية في قشور البيض وعظام الابقار تتكون اساسا من عناصر الكالسيوم والفوسفور بالاضافة لكميات قليلة من عناصر الكربون والاكسجين والصوديوم والماغنسيوم بالاضافة الي السيليكون. مع الحصول على نسبة مولارية لعنصر الكالسيوم مع الفوسفور (P/Ca) في الهيدروكسي أباتيت المحضرة معمليا مساوية 1.67 ولكن تزيد هذه النسبة بسبب تواجد كميات زائدة من كالسيوم الروابط البينية في الجينات الكالسيوم.
بالنسبة لعينات CAB و CAMK و CAD فيوجد اشعاع لعنصر السيليكون مُثبتاً مساهمة مركبات السيليكون في هذه المخاليط بنسب كبيرة والتي تساعد في التخلص من كاتيونات العناصر الثقيلة بدرجة كبيرة.
التعرف علي المجموعات الوظيفية من قياسات الاشعة تحت الحمراء بعد عملية الامتصاص (iv)
اشارت نتائج FTIR انه لايوجد تغير كبير في مواضع حزم الامتصاص بعد امتصاص ايونات المعادن وكذلك لم يتم الكشف عن اي ذٌرو امتصاص جديدة. ومع ذلك , هناك قليل جدا من الإزاحة لذروات الامتصاص , ويمكن ان يٌعزي ذلك الي اشتراك واحلال المعادن في الشبكة البلورية لألجينات الكالسيوم ومخاليطة المختلفة.
(v) قياسات حيود الاشعه السينية لعيينات للمركبات المحضرة بعد عملية الامتصاص
أجريت تحاليل الاشعة الشعة السينية XRD لعينات CA وCAHA(I) باعتبارهما الاعلي في عملية الامتصاص. تظهر النتائج انه لايوجد اطوار جديدة بعد امتصاص ايوني الرصاص الثنائي والحديد الثلاثي مما يدعم افتراض عدم حدوث الامتزاز نتيجة ميكانيكية التفكك والترسيب وانه ربما تحدث نتيجة عملية الامتزاز مثل التبادل الايوني وتكوين المتراكبات , وتظهر النتائج ايضاً حدوث تغيرات طفيفة في الشدة النسبية والحجوم البلورية وكذلك قيم التباعد d- spacing ويمكن ان يعزي ذلك الي التبادل الايوني بين ايونات المعادن و (1) ايونات الكالسيوم الموجودة في ألجينات الكالسيوم في CA و CAHA(I) كما في التفاعل التالي :
Ca(ALG)2 + Pb2+ Pb(ALG)2 + Ca2+
أو (2) ايونات الكالسيوم الخاصة بالهيدروكسي اباتيت في عينة ال CAHA(I) مكونا طور جديد من الهيدروكسي بيرومورفيت والتي لم تظهر في تحاليل XRD ربما لانها تكون بنسبة ضئيلة جدا تحت المدي الحثي XRD.
Ca10 (PO4)6 (OH)2 + x Pb2+ x Ca2+ + Ca10-xPbx (PO4)6(OH)2
(vi) التعرف على الشكل المورفولوجى والتحليل النوعي والكمى للعينات المحضرة باستخدام قياسات الميكروسكوب الإلكترونى الماسح للضوء وطاقة الأشعة السينية المشتتة بعد امتصاص ايونات الفلز.
كشفت النتائج عن بعض التغيرات في الشكل والبنية المجهرية ل CA و CAHA(I) عند التفاعل مع ايونات معادن الحديد والرصاص علاوة علي ذلك اشارت نتائج EDX الي وجود ايونات Pb2+ وFe3+ , كما تظهر النتائج أن ايونات الحديد الثلاثي هي اكثر كثافة من ايونات الرصاص الثنائي وهذا يوكد علي ان ايونات Fe3+ اكثر ازالة بواسطة CA و CAHA(I) مقارنة بايونات Pb2+. وتظهر النتائج ايضا انخفاض في نسبة ايونات الكالسيوم بعد عملية الامتصاص عنها قبل عملية الامتصاص, وهذا يمكن إيعازه الي احتمالية حدوث عملية الامتصاص نتيجة التبادل الايوني والذي يتوافق مع نتائج قياسات FTIR و تحاليل XRD. أيضا ربما تحدث عملية الامتصاص نتيجة ارتباط مخلبي لأيونات الفلزات مع مجموعتين كربوكسيليتين من الألجينات ومجموعة أواثنين من مجموعات الهيدروكسيل للألجينات , وفي هذه الحالة ربما يُكًون أيون الفلز متراكبات مع حلقتين متجاورتين لسلسلة بوليمرية واحدة من الألجينات او حلقتين لسلسلتين متوازيتين.
(vii) دراسة معدل امتصاص ايونات الفلزات تحت تاثير زمن التلامس
تمت دراسة تأثير الزمن علي قدرة ألجينات الكالسيوم ومخاليطة المختلفة في التخلص من ايونات الرصاص عند إس هيدروجيني للمحلول المحضر وهو pH=5.7 و عند 4pH = وايونات الحديد عند الاس الهيدروجيني للمحلول المحضر وهو 2.6 , وتشير الدراسة الي انه تم تحقيق نقاط الاتزان لجميع المخاليط في خلال من 10 الي 60 دقيقة من بداية الامتزاز ومراحل امتزاز (5-30) دقيقة , بنسب ازالة تتراوح ما بين66% إلي 83% للرصاص عند pH=5.7 و من 68% إلي 87% عند pH=4 وتتراوح مابين86% إلي 99% للحديد عند pH=2.6 .
وتكون عملية الامتزاز سريعة في المراحل الاولي ومتساوية تقريبا عند ازمنة التلامس العالية ويٌعزي هذا السلوك الي توافر عدد كبير من المواقع النشطة خلال المراحل الاولي من عملية الامتزاز وبالزيادة الكبيرة في زمن التلامس لم يُلاحظ حدوث أي عملية امتزاز وذلك بسبب الانخفاض في معدل الانتشار حيت ان جميع المواقع قد تم تغطيتها بايونات الفلز.
(viii) دراسة معدل امتصاص ايونات الفلزات تحت تاثير درجة الحموضة
تم دراسة تغيير درجة الإس الهيدروجيني علي عملية الامتزاز بالنسبة لايونات الرصاص فقط حيت ان ايونات الحديد Fe3+ تكون ثابته فقط عند اس هيدروجيني اقل من 3. تمت الدراسة عند قيمتين للإس الهيدروجيني وهما 5.7 (pH للمحلول المحضر) وعند pH=4, تركيز 100 ملجرام/لتر من ايونات الرصاص وكتلة من المادة المازّة ( (adsorbent 20 ملجرام/ 10 ملليتر من محلول ايون الفلز عند درجة حرارة 25 درجة مئوية. اظهرت الدراسة ان كفاءة عملية الإمتزاز عند 4 = pH تكون أعلي منها عند 5.7 = pH لجميع المخاليط.
(ix) دراسة معدل امتصاص ايونات الفلزات تحت تاثير كتلة الممتز
تم دراسة تاثير كتلة المادة المازّة بالنسبة لمعدل الامتصاص وذلك بالنسبة لإمتزاز الرصاص علي مركب ألجينات الكالسيوم (كنموذج قياسي لبقية المخاليط). تمت الدراسة باستخدام كتل مختلفة 10 و 15 و 20 ملجرام من المادة المازّة و10مليلترمن محلول الرصاص بتركيز 100 ملجرام/لتر و5.7 = pH ودرجة حرارة 25 درجة مئوية وازمنة تلامس من 5 الي 30 دقيقة. وتظهر النتائج ان معدل الامتصاص يزداد بزيادة كتلة المادة المازّة وهذ يٌعزي الي زيادة مساحة السطح ومن ثم زيادة المواقع النشطة المتاحة لعملية الامتزاز وزيادة كفاءة الازالة.
(x) دراسة كيناتيكية لعملية الإمتزاز
تمت دراسة حركية عملية الأمتزاز لمعدل امتصاص أيونات الفلزات علي سطح ألجينات الكالسيوم ومركباته المختلفة وتمت الدراسة عند ظروف تجريبية ( 25 درجة مئوية , كتلةمن المادة المازّة تساوي 20 مليجرام لكل 10 مليلتر من محاليل ايونات Pb2+ وFe3+ بتركيز ابتدائي 100 ملجرام/لتر) , وتم اختبار نموذجين حركيتين شائعتين تحت الظروف التجريبية وهما: معادلة الرتبة الأولى الكاذبة لـ Lagergren ومعائلة الرتبة الثانية الكاذبة لتحليل معدل امتزاز أبونات الفلزات علىCA و CACs.
1- نموذج الرتبة الأولى الكاذب
بتطبيق معادلة لاجرجرين لتفاعل الرتبة الاولى الزائفة التالية :
ln (qe –qt ) = ln qe – k1 t
حيث ان K1هو ثابت معدل لاجارجرين للإمتزاز (دقيقة-1) وqe وqt هى كميات العناصر الممتزة (ملجرام/جرام) عند الإتزان وعند الزمنt . ومن العلاقة البيانية بين Log(qe – qt) والزمن t, اتضح أن قيم معامل الارتباط R2تكون بعيدة وغير مقتربه من الوحدة كما ان قيم qe الناتجة من معادلة الرتبة الاولي الكاذبة تختلف تماما عن qe التجريبية , وهذا يشير الي أن معادلة Lagergren من الدرجة الأولى غير مناسبة لوصف امتزاز أيونات الفلزات المستهدفة بواسطة CA و CACs المستخدمة.
2- نموذج الرتبة الثانية الكاذب
بتطبيق معادلة تفاعل الرتبة الثانية الكاذبة التالية
t/qt = 1/k2qe2 + t/qe
على البيانات المعملية لإمتزاز ايونات الرصاص والحديد بواسطة ألجينات الكالسيوم ومركباته , حيث أن 2k هو ثابت معدل تفاعل الرتبة الثانية الزائفة (جم / ملجم. دقيقة) وqe وqt الكمية الممتزة في وحدة الكتلة عند الإتزان وعند الزمن t ، ومن العلاقة البيائية بين t/qt و الزمن t يتضح ان qe الناتجة من العلاقة الخطية لمعادلة الرتبة الثانية الكاذبة متوافقة تماما مع qe التجريبية , ووجد ايضا أن قيم معامل الارتباط تقترب جدا من الوحدة , وهذا يشير الي أن معادلة Lagergren من الدرجة الثانية مناسبة لوصف امتزاز أيونات الفلزات المستهدفة بواسطة CA و CACs المستخدمة
3 - تحديد الخطوة المتحكمة في التفاعل
من المعروف انه يوجد ثلاثة خطوات في أي عملية إمتزاز من محلول لمادة ممتزة وهي كالتالي (1) انتشار الطبقات الحدودية او الانتشار الفيلمي (2) انتشار بين الجزيئات او ثقبي (3) امتزاز علي المواقع او الاماكن الداخلية وهذه الخطوة سريعة جدا وربما لاتلاحظ , لذا فالخطوتين الاوليتين هما الخطوتين المتحكمتين في معدل الامتزاز. من اكثر الطرق المستخدمة في معرفة ميكانيكية عملية الامتزاز هي باستخدام الرسم البياني للانتشار بين الجزيئات للعالِمين ويبر وموريس بالعلاقة التالية:
qt = Kid t0.5 + C
حيث ان C ثابت و Kid هي ثابت الانتشار بين الجزيئات (ملجرام/جرام.دقييقة0.5 ( وqt هى كمية العناصر الممتزة (ملجرام/جرام) عند الإتزان. من العلاقة البيانية بين qt و t0.5 يتضح ان خطوة الانتشار بين الجزيئات ليست الخطوة المتحكمة في عملية الامتزاز لان العلاقة الخطية لا تمر بنقطة الاصل , كما توضح النتائج ان الامتصاص السريع للايونات في المرحلة الاولي يكون نتيجة الانتشار الفيلمي او تاثير الطبقات الحدودية وبعد اكتمال الأسطح الخاجية تتجه الأيونات للأنتشار بين الجزيئات اوخلال الثقوب.
(xi) دراسة الامتزاز عند درجة حرارة ثابته
من المفترض إن يكون إمتزاز ايونات العناصر + Pb2وFe3+ من الماء بواسطة CA و CAHA(I)
له سلوك يتلائم مع نموذج الإمتزاز عند ثبات الحرارة حيث أن المادة الممتزة تحافظ على الإتزان الديناميكي بين الإمتزاز وعدم الإمتزاز عند درجة حرارة ثابتة ويمكن تمثيل هذا النموذج بإستخدام معادلة لانجمير أو فريندلش. معادلة لانجمير لحساب أقصى قيمة لإمتزاز العناصر والتي تمثل بخط مستقيم تعطى من العلاقة
Ce/qe =Ce/qm + 1/ KL qm
حيثCe هو تركيز الإتزان للعنصر الثقيل المتبقي في المحلول(ملجرام/ لتر) عندما تمتز منه كمية تساوي qe. و qe هى الكمية الممتزة عند الإتزان ( ملجرام/ جم) وqm هى سعة الإمتزاز القصوى التي ترجع الى حدوث تغطية كاملة للسطح من طبقة واحدة (ملجرام / جم) و kLهوثابت لانجمير الذي يتناسب عكسيا مع طاقة الإمتزاز (ملجرام) ويمكن حساب qe من المعادلة
qe =((Co-Ce)V)/m
حيثCo هو تركيز الأيون الإبتدائي (ملجرام/ لتر) و Ct التركيز النهائي للعنصر (ملجرام/ لتر بعد مرور فترة من الزمن t , و V هو حجم المحلول الابتدائي (لتر) و m هي كمية الجينات الكالسيوم او مركبه مع الهيدروكسي اباتيت المضافة , كذلك من العلاقة التالية :
RL=1\1+ KLCe
نستطيع معرفة ما اذا كانت عملية الامتزاز مفضلة او غير مفضلة او غير عكسية او خطية حيث ان RL هي معيار الاتزان بلا ابعاد فاذا كانت قيمة RL تساوي صفر فان عملية الامتزاز تكون غير عكسية , واذا كانت تساوي 1 فان الامتزاز يكون خطي , واذا كانت بين صفر و واحد تكون مفضلة واذا كانت اكبر من الواحد تكون غير مفضلة.
اما معادلة فريندلش والتي تفترض وجود سطح غير متجانس مع توزيع غير منتظم لحرارة الامتزاز علي السطح وان عملية الامتزاز تتم علي طبقات عديدة و تمثل تلك المعادلة بخط مستقيم و تعطي من العلاقة
log qe =log KF + (1/n) log Ce
حيث KF و n هما ثوابت فريندلش ومن خلال قيمة (1/n) نستطيع معرفة ما اذا كان الامتزاز غير انعكاسي اذا كانت تساوي صفر اما اذا كانت بين 0 و 1 فان عملية الامتزاز تكون مفضلة اما اذا كانت قيمة (1/n) اكبر من 1 تكون عملية الامتزاز غير مفضلة.
من تطبيق تلك المعادلات علي النتائج التي تم الحصول عليها من عملية الامتزاز يتضح ان امتزاز ايونات الرصاص علي الجينات الكالسيوم يمكن تمثيلها بنموزج لانجمير بمعامل ارتباط جيد مايعني تغطية ايونات الرصاص لطبقه واحدة من CA , كذلك قيمة RL=0 تشير الي ان عملية الامتزاز تكون غير انعكاسية.
اما في حالة إمتزاز ايونات الرصاص والحديد علي سطح CAHA(1) وكذلك امتزاز ايونات الحديد علي سطح CA فيتضح أنها تُمثل جيدا بنموذج فريندلش بمعامل ارتباط جيد وكذلك قيمة (1/n) تشير الي ان عملية الامتزاز تكون مفضلة.
من خلال كل ماسبق من نتائج التحاليل المختلفة ومن دراسة كيناتيكية عملية الامتزاز يتضح أن:
إزالة عناصر الحديد والرصاص ربما تتم عن طريق تكوين متراكبات علي سطح ألجينات الكالسيوم ومركباته المختلفة او تتم عن طريق التبادل الايوني بين ايونات تلك العناصر وايونات الكالسيوم في عينات ألجينات الكالسيوم ومركباته المختلفة او ايونات الكالسيوم المرتبطة بالهيدروكسي اباتيت في عينات ألجينات الكالسيوم التي تحتوي علي مادة الهيدروكسي اباتيت.
مركب CAG هوالأعلي كفاءة في أزالة عنصر الرصاصPb2+ يليه مركبات CA و CAHA(I) ومركب CAHA(I) هو الأعلي كفاءة في إزالة عنصر الحديد Fe3+ يليه CAHA(II) ثم CA.
}
رسالة مقدمة من الطالب
محمد أحمد عبده عبدالله الامير
بكالريوس علوم (كيمياء – 2009)
كجزء من متطلبات الحصول علي
درجة الماجستير في العلوم
(كيمياء)
(كيمياء غير عضوية وتحليلية)
إلى
قسم الكيمياء
كلية العلوم
جامعة قناة السويس
الإسماعيلية
(2019)
دراسات علي إزالة العناصر الثقيله من المياه الملوثه باستخدام بعض المواد العضوية وغير العضويه المركبه
لجنة الأشراف التوقيع
1- أستاذ دكتور / صبري عبد الحميد القرشي ..........................
أستاذ الكيمياء غير العضوية
قسم الكيمياء
كلية العلوم
جامعة قناة السويس
2- دكتور / أيمن عبد المؤمن محمد مصطفي ..........................
مدرس الكيمياء الفيزيائية
قسم الكيمياء
كلية العلوم
جامعة قناة السويس
3- دكتور/ عباس ممدوح عباس ..........................
مدرس الكيمياء غير العضوية والتحليلية
قسم الكيمياء
كلية العلوم
جامعة قناة السويس
وافق مجلس الكلية بتاريخ / /
كما وافق السيد الأستاذ الدكتور / نائب رئيس الجامعة بتاريخ / /
علي منح درجة الماجستيرفي العلوم للطالب
محمد أحمد عبده عبدالله الامير
عنوان ارسالة:
دراسات علي إزالة العناصر الثقيله من المياه الملوثه باستخدام بعض المواد العضوية وغير العضويه المركبه
لجنة الحكم والمناقشة التوقيع
1- أستاذ دكتور / صبري عبد الحميد القرشي ..........................
أستاذ الكيمياء غير العضوية
قسم الكيمياء
كلية العلوم
جامعة قناة السويس
2- أستاذ دكتور / عصام عبدالعزيز ابراهيم كيشار ...........................
أستاذ الكيمياء غير العضوية
قسم الكيمياء
بنات عين شمس
3- أستاذ دكتور مساعد / خلود محمد ابو النور ...........................
أستاذ مساعد الكيمياء التحليلية
قسم الكيمياء
كلية العلوم
جامعة قناة السويس
وكيل الكلية لشئون الدراسات العليا عميد الكلية
أ.د/ علاء الدين عبدالعزيز سلام أ.د/ محمد سعد زغلول
.