While MOFs based on divalent metals, such as Zn2+ and Cu2+, have received much attention over the past decade, less progress has been achieved on the synthesis of new MOFs containing tri- and tetravalent metals [22]. Among these, vanadium MOFs are particularly rare. Applications of vanadium-based-MOFs can be found in areas of chemistry such as catalysis, adsorption, separation of N2, CO2, CH4, magnetism, and others [23–25]. Ferey and coworkers reported three-dimensional vanadium (III) dicarboxylate, derived from terephthalic acid in HF, MIL-71. MIL-71 was prepared from a mixture of metallic vanadium, HF, 1,4-benzenedicarboxylic acid, and deionized water heated 3 days at 473 K under hydrothermal conditions. MIL-71 exhibits two features according to the authors: (i), it is the first solid with a two-dimensional inorganic subnetwork among the series of hybrid vanadocarboxylates, and (ii) compared with other trivalent cations, it exemplifies once more the peculiar behavior of V(III), which easily oxidizes into V(IV) [26]. Finally, the Vanadium atoms show an octahedral geometry where four oxygen atoms coordinated to the metal occupy the equatorial positions and two fluorine atoms that occupy the apical positions as observed from **Figure 9**.

More interesting to observe is the pore size obtained in MIL-71 in the different planes. **Figure 9** represents the pores in the MOF, the left side denotes the 010 planes with distances of 3.57 Å × 10.72 Å and the right plane 100 with distances of 3.57 Å × 3.87 Å.

*Transition Metals-Based Metal-Organic Frameworks, Synthesis, and Environmental Applications DOI: http://dx.doi.org/10.5772/intechopen.104294*

#### **Figure 9.**

*Molecular representation, where the vanadium atom is hexacoordinated displaying an octahedral geometry.*

The right side represents the front part, and the left side represents the lateral view, forming in this way a rectangle. Red color represents oxygen atoms, green color represents fluorine atoms, gray color represents carbon atoms, and purple color represents vanadium atoms. Metals and many small molecules could be adsorbed in this material due to the great number of electrons spported by the oxygen and fluorine atoms and the aromatic rings in the MOF, **Figure 10**.

#### **4.4 Chromium MOFs**

Chromium-based MOFs have been widely used with different applications such as environmental remediation, gases sorption, catalysis, carrier systems, among others [27–31]. Feng and coworkers have recently reported an ultrastable High-Connected Chromium Metal-Organic Framework, which was obtained using a Teflon cup, weighing, and combining chromium nitrate nonahydrate, 1,4-benzenedicarboxylic acid, 2,4,6-tri(4-pyridyl)-1,3,5-triazine, and 50 μL hydrofluoric acid dissolved in water. The solvothermal technique indicates that after being stirred for an hour, the vessel was sealed and was subsequently placed in a 220°C preheated oven for 2 days.

**Figure 11.** *Illustration showing graphic depicts of the porosity of Cr-based MOF.*

As mentioned before, the importance to follow the technique is imperative to obtain the correct product, because a minimal variation of temperature, reaction time, pH, and even stir velocity can conduce to a different product. The obtained MOF gives rise to a tridimensional structure, and the material presents high porosity. The chromium atoms in the structure are hexacoordinated showing an octahedral geometry with typical bond distances in the range from 1.94 Å to 2.17 Å, **Figure 11** [32].

#### **4.5 Manganese MOFs**

Manganese-based MOFs are very stables, and the reports in the literature are very large, which include catalysis, materials with magnetic properties, transport, energy, gas sorption, among others [33–37]. Manganese(II) ions/clusters have been used as secondary building units (SUBs) to build MOFs in fields of gas adsorption and magnetism attributed to their specific electronic configuration.

Wang and coworkers reported a manganese-based MOF applied in the decomposition of ozone, which means that the MOF works like a catalytic species to decompose the ozone in water. The technique employed for the synthesis was solvothermal, weighing the corresponding moles of Mn (NO3)2·6H2O and the linker H4TTPE, which were dissolved in a mixture of DMAc/H2O. The yellow solution obtained was stirred at room temperature for 5 min, sealed in a 25 mL Teflon-lined bomb, and heated at 120°C for 72 h [38].

The material shows according to the X-ray structure, two different kinds of manganese atoms, one manganese being pentacoordinate with a very distorted trigonal bipyramidal geometry, and the second manganese being surrounded by six atoms accommodated in an octahedral fashion. The X-ray shows two different types of pores in the structure, and the water molecules are coordinated to the manganese metal atom. The two water molecules occupy part of the interior of the pores, and although both pores are similar in size, it could be a slight difference in sorption capacity when small molecules go inside the pore and interact either with the benzene electron-rich rings pores or the more exposed electron-rich pores that contain oxygen and nitrogen atoms from the tetrazole moiety, **Figure 12**.

*Transition Metals-Based Metal-Organic Frameworks, Synthesis, and Environmental Applications DOI: http://dx.doi.org/10.5772/intechopen.104294*

**Figure 12.** *Manganese-based MOF shows the pores for possible sorption of ions and molecules.*

#### **4.6 Iron MOFs**

Iron-based MOFs have been widely reported in biological applications due to their high stability, ease of synthesis, and low toxicity of the metal to the human being. Application in areas such as environmental remediation, adsorption of volatile compounds, catalysis, drug delivery, water remedying, glucose biosensing, among others [39–43].

Long and coworkers described an iron MOF that was synthesized from anhydrous ferrous, 1,4-dihydroxyterephthalic acid, DMF, and methanol. The reaction mixture was heated at 393 K and stirred for 18 h to afford a red-orange precipitate. The solid was collected by filtration and washed with 100 mL of DMF to yield 2.0 g (91%) [44].

Although the reaction implies a solvothermal method for 18 h, worth it for the high yielding obtained. In **Figure 12**, we observe a small representation of the MOF, orange, gray, and red spheres represent Fe, C, and O atoms, respectively, and the hydrogen atoms were omitted for clarity. The iron atom possesses two different coordination ways, square pyramidal and octahedral. The pores are well defined and the oxygen atoms possessing two-electron lone pairs can interact with different metals and molecules. Even more, all the channels are well defined, and **Figure 13** shows a small portion of the crystal.

#### **4.7 Cobalt MOFs**

Cobalt-based MOFs have been broadly studied due that cobalt salts being cheap and easy to obtain, even more, the cobalt atom can form Penta or Hexa coordinate geometries, increasing the possibilities of coordination modes. These materials found applications in oxygen and hydrogen evolution, magnetism and superconductivity, catalysis, electrocatalysis, synthesis of nanomaterials, and more [45–49].

Burgos and coworkers could develop nanosheets of cobalt MOF for enhanced electrocatalytic water oxidation. The X-ray studies show that there exists just one type of

cobalt atom, which is hexacoordinate with an octahedral geometry, and the pyridine solvent is coordinated to the cobalt atom. The pyridine rings are close enough to interact via π-π stacking with an average distance of 9.25 Å. The cobalt atoms are hexacoordinated showing a distorted octahedral geometry with coordination distances in the range of 2055–2179 Å**.** Because this material contains two different types of cavities, both with an electron-rich environment, metals or small molecules can interact in the surroundings to produce stable states [50].

The use of a well-defined cobalt cluster as the starting compound for the synthesis directs the construction of a Co-MOF with an unusual topology. In this MOF, the layered double nanosheets are held together by π − π stacking interactions between labile pyridine ligands. It has been shown that this material delaminates in the presence of water and that the original 3D layered structure can be regenerated by solvothermal treatment with pyridine so that the individual nanosheets have associated memory **Figure 14.**

#### **4.8 Nickel MOFs**

The combination of porosity and the presence of coordinatively unsaturated Ni2+ sites are also of special interest because of catalytic properties and the strong H2 binding affinity. Added to this, the chemical and thermal stability and the presence of accessible Lewis acid sites are some of the reasons why Nickel-based MOFS find application in a great variety of fields of chemistry and reports can be found on catalysis for the ethylene oligomerization, CO2/CH4 separation, magnetic and conducting materials, methanol oxidation, among others [51–55].

Gong and coworkers reported the synthesis of Nickel-based MOFS using the microwave-assisted [MA] technique combined with the solvothermal reaction. In contrast to conventional methods, MA methods allow the rapid and systematic investigation of large synthesis fields. This enables the effective discovery of new

*Transition Metals-Based Metal-Organic Frameworks, Synthesis, and Environmental Applications DOI: http://dx.doi.org/10.5772/intechopen.104294*

**Figure 14.** *A small part of the crystal shows the proximity of the pyridine groups that originate the π-π interactions.*

compounds, the fast optimization of synthesis conditions, and because of the large amount of data, it allows the extraction of reaction trends. MA methods have been successfully applied in the investigation of porous MOFs.

The resulting material was obtained as a crystalline solid and the X-ray diffraction showed the structure that possesses a hexacoordinated cobalt atom arranged in an octahedral fashion with a Ni-Ni interaction of 2.713 Å. The coordination bond distances are in the range from 1.972 to 2.009 Å [56]. Part of the structure is shown in **Figure 15**.

#### **4.9 Copper MOFs**

Reports on MOFs containing copper include novel application in coordination strategies to control the growth orientation of the crystals, modification, and adsorption of different volatile organic compounds, MOF nanoparticles applied for sensitive fluorescent detection of ferric ion, copper-based MOFs for sensitive colorimetric detection of staphylococcus aureus, separation of CO2 over CH4 or N2, among others [57–61].

Park and coworkers adopted a facile green synthesis for the preparation of a copperbased MOF applied in the cycloaddition reaction of CO2 and epoxide. The coordination geometry around each Cu in the bimetallic cluster is octahedral, where the square base is established by the four BDC units with their four carboxylate oxygen atoms. The PNU-25 structure contains 3D channels with rectangular windows of dimensions 15.21 × 10.80, which undergo interpenetration *via* various supramolecular interactions forming an overall triple interpenetrated network, as depicted in **Figure 16** [62].

The heterogeneous PNU catalysts efficiently catalyzed the synthesis of cyclic carbonates by the coupling of epoxide and CO2 under ambient pressure and lower reaction temperature. The PNU catalysts demonstrated remarkably good thermal stability for the cycloaddition reaction. The coordinatively unsaturated Cu(II) units

**Figure 15.** *Part of the structure of Ni-based MOF obtained by microwave-assisted technique.*

#### **Figure 16.**

*The PNU-25 structure contains 3D channels left, with copper metal atoms being hexacoordinated within an octahedral geometry, right.*

and the basic N atoms resulted in a large number of acidic-basic sites, facilitating the conversion of epoxides.

#### **4.10 Zinc MOFs**

One of the important applications of zinc-based MOFs is the separation of materials, gases, and compounds. By using the modified MOFs, diverse gases, organic and inorganic compounds were separated, i.e., H2/CO2, Xe, and Kr, O2/N2/CH4, other applications such as catalysis in organic synthesis, removal and detection of antibiotics in water, sensing, photocatalytic activity, semiconductive and magnetic properties, thin-film nanocomposite membrane incorporated with Porous Zn-Based Metal-Organic Frameworks, among others [63–67].

*Transition Metals-Based Metal-Organic Frameworks, Synthesis, and Environmental Applications DOI: http://dx.doi.org/10.5772/intechopen.104294*

**Figure 17.** *A small fraction of Zn-ttb-bdc MOF shows two coordination water molecules.*

Yang and coworkers develop a low toxicity MOF for the detection of organic and inorganic contaminants from water. The material was obtained according to the mixture of the linker Httb and Zn (NO3)2·6H2O in a mixture of solvents, H2O, and CH3CN, and the mixture was stirred for 20 min before adding in a steel vessel at 160°C for 5 days to obtain the Zn-based MOF, with yield 64%.

The MOF exhibits high water and chemical stability as well as excellent fluorescence properties. The remaining binding sites show higher sensitivity and better fluorescence response to the representative organic micropollutant TNP and inorganic pollutants (Fe3+ and Cr2O7 2− in wastewater. The Zn metal is hexacoordinated to four nitrogen atoms from the tetrazole ring and two oxygen atoms belonging to water molecules, **Figure 17** [68]. The final structure presents π-π interactions between the aromatic rings.

#### **5. Synthesis OF MOFs**

There are many reports in the literature about the ways and techniques commonly used, frequently called reaction conditions, which are in some cases very specific to synthesize the metal transition-based MOFs. The different synthesis methods that have been applied in the last 20 years in addition to room temperature synthesis are conventional electric (CE) heating, microwave (MW) heating, electrochemistry (EC), mechanochemistry (MC), and ultrasonic (US) methods that have been employed.

Conventional step-by-step methods, as well as high-throughput methods, have been employed in some of the studies. In addition to the crystal size or shape, thin films, membranes, and various other shapes made of MOFs have been reported, which require the application of different and specific synthesis methods. Nevertheless, many other synthetic methods and parameters, such as temperature, reaction time, pressure, pH, and solvent, must be considered as well. Numerous different synthetic approaches, including slow diffusion, hydrothermal, and solvothermal, can be applied to produce MOFs relying on the resulting structures and features [69].

**Figure 18.** *Strategies and methods for the synthesis of crystalline materials [2].*

**Figure 18** presents a resume of the different techniques applied for the synthesis of MOFs, where we observe the conventional heating, sonochemistry, microwaveassisted, using different temperature programs that let the correct nucleation and growth of single crystals.

#### **6. Environmental applications of transition metal-based MOFs**

As could be discussed in the previous part, there are many different techniques to achieve the synthesis of MOFs, and something very important to consider is the combination of procedures that can allow the formation of materials with specific pore sizes and shapes. The environmental applications of MOFs depend on the final structure obtained, which is very related to the technique employed and the metals and linkers used for the synthesis. Pore size and electron charge distribution into the pores and the surface let those metals or even neutral or charged molecules interact with each other. In this section, we'll focus on three environmental applications, heavy metal atoms, and azo dyes removal from water, hydrogen evolution, and catalysis directed to the transesterification process of non-edible oils to obtain biofuels.

#### **6.1 Metals and azo dyes removal from wastewater**

There are many reports in the literature of heavy metal atoms removed from aqueous mediums using crystalline materials. The remotion of those metals from wasted water is due to the weak interactions generated between the metals and the electron-rich environment of the pores. In the same way, the remotion of azo dyes

#### *Transition Metals-Based Metal-Organic Frameworks, Synthesis, and Environmental Applications DOI: http://dx.doi.org/10.5772/intechopen.104294*

(medium-size molecules). i.e., Congo red, methylene blue, and methyl orange from wastewater, has been widely described, due to those azo dyes are considered toxic and carcinogenic pullulans in the water.

Heavy metals such as arsenic, cadmium, chromium, copper, lead, mercury, nickel, and zing with a density over 5 g cm−3polluting our water are a rapidly growing global concern. These elements can be found within the environment, be it in water reservoirs, the atmosphere, or soil [70]. The recovery and elimination of toxic metal ions from wastewater are of concern with increasing awareness toward the need for protecting nature. Mercury and lead are the most toxic species and can cause bioaccumulation in kidneys, brain, lung tissues, gastrointestinal tract, central nervous system, and reproductive system [71].

To date, several approaches have been applied to remove the heavy metals from water including nanofiltration, membrane separation, ion exchange, resin, photocatalytic degradation, chemical moisture, membrane filtration, freezing, chemical deposition, biological treatment, reverse osmosis, adsorption, etc. [72].

Studies to investigate the sorption mechanisms governing the sorption process and to determine whether the sorption mechanism are controlled by a chemical or physical mechanism, Morsali et al. applied conventional sorption kinetics models such as pseudo-first-order, pseudo-second-order, and intraparticle diffusion to the study. The correlation coefficient of the pseudo-second-order model for Cd(II) is close to 1 (R2 = 0.9999) and exhibited that the pseudo-second-order model is more consistent with experimental data than the other model. Therefore, it is compatible with chemical sorption, which indicates that transfer, exchange, or sharing of electrons has taken place. Therefore, supramolecular interactions between cadmium ion and the free electrons of ligand nitrogens and the electrostatic interaction between the cation and the dipole of nitrogen on the dihydropyrazine ring of ligand are proposed as the dominant interactions between the metal ion and the structure [73].

**Figure 19.** *A small portion in MOF TMU-60 and pore size obtained after the synthesis.*

The X-ray structure showed that the aromatic rings from the linker are available to interact through supramolecular interactions with different metals or even small molecules through week interactions **Figure 19**.

The synthesis was achieved in DMF at 60°C, and then the solution was transferred to a Teflon autoclave and heated to 120°C for 72.

Large amounts of dyes are commonly used in many manufacturing, such as textiles, clothing, and printing. Among these, azo dyes are one of the largest groups that are heavily used, but most difficult to be degraded, and even more, the azo dyes and their derivative products are toxic to the aquatic environment and are mutagenic and carcinogenic to humans. Therefore, the treatment of wastewaters containing these dyes is necessary [74].

Numerous studies have considered new processes to eliminate emerging organic contaminants (EOCs) from water, i.e., ozonization, chlorination, sonodegradation, biodegradation, inorganic heterogeneous catalysis, activated carbon treatment, and more. Among these technologies, metal-organic frameworks (MOFs) have been recently investigated for the removal of contaminants in water [75].

Separation or degradation requires strong research efforts to modify MOFs via controlling their pore diameter (adding functional groups, creating defects) or to construct MOF composites to engineer materials as improved adsorbents and/or catalysts for contaminated water treatment. MOFs and MOF composites have been proposed for the removal of a wide range of contaminants, including dyes, pharmaceuticals, plasticizers, herbicides and pesticides, industrial products, among others. While a few reviews have already documented the use of MOFs in the removal of novel contaminants [76].

Choe et al. reported the adsorptive removal of various dyes using a Zr porphyrinic MOF, PCN-224. The plausible mechanism for adsorptive removal revealed multiple interactions between the dye and porphyrin linker/Zr6 node via π − π interactions and hydrogen bonding, respectively. Such results demonstrate that PCN-224 is an excellent adsorbent, providing superior water stability, pore aperture of suitable size, and multiple interaction sites **Figure 20** [77].

The best performance of dye adsorption onto PCN-224 comes from structural properties such as appropriate pore aperture, volume as well as various types of

interactions such as π−π interaction, hydrogen bonding, and electrostatic interaction. Significantly, an interaction between the sulfonate group of the MO molecule and the Zr6 node of PCN-224 was demonstrated through experimental and theoretical studies. Arrangement of tripod form between Zr node and sulfonate group of the MO molecule can form hydrogen bonding.

#### **6.2 Hydrogen evolution**

Highly efficient hydrogen evolution reactions (HERs) will determine the mass distributions of hydrogen-powered clean technologies in the future. That's why hydrogen evolution is one of the topics more explored in the last years because molecular hydrogen is the best environmentally friendly fuel available, which reacts with oxygen to produce energy and water.

The procedures employed to produce hydrogen are still expensive because of the catalysts employed to obtain the gas and the conditions and hence the high-cost obtention. Current industrial hydrogen production methods include coal gasification (followed by water-gas shift reaction), steam reforming, cryogenic distillation, and water splitting.

The specific surface area of MOFs can range from 1000 to over 6000 m2 g−1, thanks to their tailorable porous structures, which play a significant role in enhancing the catalytic HER process. Large surface area and pore volume ensure sufficient contact between the electrolyte (or reactant solution) and the surface of the catalyst, which essentially improves catalytic performance by exposing more active sites for the catalytic reactions to take place. It is that research on MOFs for HER primarily focuses on the following three techniques: electrocatalytic, photocatalytic, and chemocatalytic HER [78].

MOFs, MOF supports, and MOF derivatives can be utilized as catalysts in the abovementioned hydrogen production methods. Most of the materials used for photocatalytic hydrogen evolution PHE applications include inorganic oxides such as TiO2, ZnO, and SrZrO3, due to their high stability. The application of metalorganic frameworks in these processes is limited due to the loss of efficiency attributed to the recombination of electrons and holes [79]. Even so, together with the unique porous structure of MOFs, a remarkable hydrogen evolution reaction HER performance can be achieved using different overpotential in phosphate buffer solution (PBS, pH = 7.0) [80].

In the same way, the significant effect of crystallinity in the photocatalytic activity of metal-organic frameworks was demonstrated through the evaluation of different samples with different crystallinity in the HER reactions. The samples with high crystallinity produce too a higher amount of hydrogen, which is attributed to the lower recombination supported by the experiments of photoluminescence and electrochemical impedance directly related to a high-ordered material.

Rivera et al. described the synthesis of a BDC-Zn MOF, which was firstly used for methyl orange and methyl blue sorption and followed by PHE under solar light. MO presented the best adsorption result, with a maximum adsorption capacity of 2100 mg/g, which is higher than all the MOFs reported in the literature. For HER, the activity was enhanced 24 times in photocatalyst with MO adsorbed, and 27 times for the MB adsorbed (from 47 to 1148 and 1259 μmol/gh, respectively). This result is attributed to better light adsorption and a decrease in charge recombination. It's important to mention that even though the reflux method presented the disadvantage that it is not possible to obtain single crystals, the reaction conditions such as

**Figure 21.** *Mechanism of HER by using the BDC-Zn-MOF obtained under solvothermal conditions [81].*

temperature, pressure, and time are more ecological in comparison with the traditional MOF 5 solvothermal synthesis and analogous [81].

**Figure 21** shows the mechanism for HER. In the first step, MO/MB molecules linker to BDC-Zn MOF captures the light from the solar simulator. Light produces the transference of electrons from the conduction (LUMO orbital analog) to the valence band (HOMO orbital analog), this produces a hole formed in the conduction band, which oxidizes water molecules, finally, the electron in the valence band reduces H+ to H2. At the same time the species O2 − and OH− are formed, and these could affect MOFs by redox reactions; however, possibly MO/MB acts as a sacrificial agent avoiding the BDC-Zn MOF structure collapse by the capture of these species. Therefore, growing hydrogen production is observed.

#### **6.3 Catalytic transesterification reaction**

Biodiesel is green and renewable energy, which is a promising substitute to replace fossil combustibles. Normally, biodiesel is produced via transesterification/esterification with the assistance of a homogeneous or heterogeneous catalyst. Biodiesel is a series of fatty acid alkyl ester (FAAE), which is normally derived from transesterification of vegetable oil or animal fat (triglyceride) and alcohol with the assistance of catalyst as shown in **Figure 22**.

In general, transesterification is the most used technique for biodiesel production, which is proceeded in three steps. In the first step, triglyceride reacts with alcohol generating monomolecular FAAE and diglyceride. Then, diglyceride reacts with alcohol resulting in monomolecular FAAE and monoglyceride. Finally, monoglyceride reacts with alcohols giving rise to monomolecular FAAE and glycerol [82].

Industrial biodiesel production is based on the transesterification of triglycerides with methanol or ethanol using stoichiometric amounts of strong Brönsted base (e.g., sodium or potassium hydroxide/methoxide) and homogeneous acid catalysis, such as H2SO4, HCl, H3PO4, but the problem with this homogeneous catalyst is their difficult recovery from the media. The ease of heterogeneous catalyst to be recovered at the

*Transition Metals-Based Metal-Organic Frameworks, Synthesis, and Environmental Applications DOI: http://dx.doi.org/10.5772/intechopen.104294*

**Figure 22.** *Transesterification reaction of triglycerides in alcohol as a solvent.*

end of the reaction just by a simple filtration reduces dramatically the cost in the biodiesel obtention process.

One of the properties of MOFs compared with other porous solids is the facile introduction of the desired active sites. Solvent/water molecules weakly bounded around metal nodes in MOFs could be removed upon activation; thus, these coordinatively unsaturated metal sites behave as Lewis acid sites [83].

MOFs can be used in combination with other materials to form composites or as catalytic support to solve two specific problems, stability, and recyclability issues. The porous materials, metal-organic frameworks (MOFs) are suitable for enzyme immobilization owing to their ideal features of tunable pore size, and topological and compositional versatility. Currently, multiple enzyme-MOF composites have been successfully created using physical adsorption or covalent attachment strategies. Also, because MOFS catalyze many reactions, they accelerate the velocity of the reaction and low the transition states of the molecules to react faster and slower reaction times and energy consumption [84].

To increase the catalytic activity in MOFs, three different strategies can be used. The first strategy is post modification of MOFs, which results in the linker modification to increase the active sites and hence the catalytic activity. The second strategy is the formation of composites, which results in mixing the MOFs with different basic or acid materials together to obtain better results and structural stability in MOFs. The third strategy is to develop new linkers that incorporate the specific active sites that remain free once the MOFs are obtained [85]. Finally, because the transesterification process requires very large reactions times, the energy consumption is high; therefore, different synthesis techniques can be used in combination. Rivera et al. could obtain through hydrothermal synthesis a 3D-Co-based MOF. The Ultrasonic-assisted synthesis is a powerful technique, the yields obtained are increased, and the reaction times are decreased using milder conditions.

Rivera et al. reported the green synthetic approach that involves the ultrasonic technique that allows in terms of energy conservation higher purity products. The high catalytic activity exhibited by the catalyst reported has to do in part with the free carboxylic acid groups that remain free in the structure, as revealed by the X-ray diffraction. The application of this catalyst resulted in 80% of total conversion after 12 h at just 60°C [86].

The use of composite MOFs is widely described in the literature, and the idea to produce biodiesel is to add new basicity or acidity sites to the material to increase the catalytic activity or simply to get better stability. Zhang et al. synthesized the reusable and highly active Fe-BTC and UiO-66 metal-organic framework by hydrothermal method, and the MOF was applied for the acid-catalyzed esterification of oleic acid with methanol. Typically, the esterification reactions of oleic acid with methanol

**Figure 23.** *Possible reaction mechanism for the esterification reaction process [87].*

were performed in a 50 mL stainless steel reactor equipped with a magnetic stirrer. The mechanism proposed for the esterification reaction is shown in **Figure 23**.

Characterization analyses indicated that the ZrSiW/UiO-66 possesses an appropriate structure and high acidity. The highest oleic acid conversion of 98.0% was obtained using the ZrSiW/UiO-66 nano-hybrids nanocatalyst under the optimal esterification reactions: 150°C, 0.24 g catalyst, 1:20 molar ratio of oleic acid to methanol, and a reaction time of 4 h [87]. The easy preparation of MOF composites lets these materials find better opportunities to be investigated [87].

#### **7. Conclusions**

Metal-organic materials are easy to synthesize and have a wide variety of applications. MOF materials are constructed from linkers and different metal atoms, which can interact via coordination bonds to produce 1D, 2D, and 3D porous structures. The final porous structure depends primarily on the linkers and metals used for the synthesis but equally important to obtain specific crystalline structures and hence high-ordered and porous materials is to apply the correct technique that involves the correct pressure, reaction time, temperature, pH value, solvent, or a mixture of solvents, stirring, or even microwave and ultrasonic-assisted synthesis, which can be applied and combined with other techniques to obtain desirable size and porous structures in the MOFs. The environmental applications of MOFs such as sorption of metals and azo dyes depend on the final structure obtained, the pore size, and the electron charge distribution into the pores, and the surface allows metals or even neutral or charged molecules to interact with the MOF to form different contacts such as hydrogen bonding, π-π interactions, or simply interact through Van der walls forces. On the other hand, the catalytic activity for the transesterification process to obtain biodiesel depends mainly on the availability of metal centers, which are acid sites that *Transition Metals-Based Metal-Organic Frameworks, Synthesis, and Environmental Applications DOI: http://dx.doi.org/10.5772/intechopen.104294*

can be increased by adding more acid sites by a post-modification synthesis or simply by using a specific linker containing the acid sites in the original structure. Finally, the hydrogen evolution reaction depends mainly on the structure of MOFs (crystallinity) and the linkers, that is, delocalized systems that allow HOMO and LUMO orbitals to get less separated, and hence the electrons can get transferred more efficiently to oxidase water.

#### **Acknowledgements**

Thanks to committee pro-mejoras from Universidad Veracruzana, Facultad de Ciencias Químicas, Orizaba, Ver. México, who supported the funding of this project.

#### **Conflict of interest**

"The authors declare no conflict of interest."

#### **Author details**

Lidia E. Chiñas-Rojas1,2, Guadalupe Vivar-Vera1 , Yafeth F. Cruz-Martínez1 , Seth Limón Colohua1 , José María Rivera1 and Eric Houbron1 \*

1 Facultad de Ciencias Químicas, Universidad Veracruzana, Orizaba, Veracruz, Mexico

2 Universidad Martí, Poza Rica, Veracruz, Mexico

\*Address all correspondence to<EMAIL_ADDRESS>
© 2022 The Author(s). Licensee IntechOpen. This chapter is distributed under the terms of the Creative Commons Attribution License (http://creativecommons.org/licenses/by/3.0), which permits unrestricted use, distribution, and reproduction in any medium, provided the original work is properly cited.

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#### **Chapter 5**

## Adsorption and Its Applications: Using Zinc Adsorption on Water Hyacinth to Elaborate the Kinetics and Thermodynamics of Adsorption

*Joseph Munene Mwaniki*

#### **Abstract**

Adsorption is a technique for removing adsorbate from the liquid or gas phase using adsorbents. The adsorbent is solid while the adsorbate can either be dissolved in liquid or gas. Adsorption has attracted the attention of many researchers because of its wide applicability in water and air purification, environment friendly, effectiveness, and ease to design as compared with the other methods. Activated carbon has been used as an effective adsorbent. However, its application is limited since it's expensive. This has necessitated research interest in other materials that are safe and economical instead of commercial activated carbon. Some of the materials that have been successfully tested include sawdust, silica gel, zeolites, clay minerals and oxides, nanomaterial, agricultural by-products, biological waste, ion exchange resins and water hyacinth, etc. Although some of these materials are effective, they are not readily available. The kinetics of adsorption is done through testing the adsorption data against standard kinetic models and the model with the best line of fit, based on the values of coefficient of determination (R2 ) is selected. The adsorption process is described using isotherms such as Freundlich and Langmuir. This chapter sheds more light on adsorption, the most common adsorbents, kinetic models, isotherms, and adsorption applicability.

**Keywords:** adsorption, adsorbents, kinetics, mechanism, isotherms, thermodynamics

#### **1. Introduction**

Over the last decade, there has been a tremendous increase in industries worldwide. Wastewater released from the industries is contaminated by toxic substances such as heavy metals. Heavy metals are persistent, fairly soluble in water, and hence easily absorbed into living cells [1]. Heavy metal pollution poses health problems such as damage to the liver, kidney, circulatory and nervous system, dermatitis, insomnia, tumor formation, rheumatoid arthritis and respiratory cancer. Heavy metals of great concern are mercury, lead, zinc, nickel, cobalt, chromium, copper, and cadmium.

Due to these and many more adverse effects of heavy metals, on humanity and the environment, there is a need for protection and restoration of the environment through the removal of heavy metals from the industrial wastewater before being released into the environment.

Several methods have been used to remove heavy metals from industrial wastewater, which include the use of membranes such as reverse osmosis, electrochemical techniques such as electrolytic extraction and electrodialysis, ion exchange, and chemical precipitation. Although these methods are effective, they are costly in terms of infrastructure, control systems, and energy [2]. There exist cheaper methods of eliminating heavy metals from wastewater such as adsorption [3]. Adsorption is becoming a regular method for the removal of heavy metals from wastewater since its relatively less costly, energy-efficient, environmentally friendly, and inexpensive [4]. Another key advantage of adsorption is the ease of designing and operating [5]. Adsorption is very effective even when the concentration of heavy metal is as low as 1 mg/L, offering metal recovery, metal selectivity and regeneration of the adsorbent material [6].

#### **2. Adsorption mechanism**

Adsorption takes place in four definitive stages: Stage 1: The mass transfer of the adsorbate by diffusion from the bulk fluid to the solid–liquid boundary layer surrounding the adsorbent particle. Stage 2: External diffusion. The transportation of the adsorbate through diffusion within the boundary layer to the outer surface of the adsorbent. Stage 3: Internal diffusion. The transfer of the adsorbate from the outer surface of the adsorbent to its inner surface by diffusion. Stage 4: The adsorption of the adsorbate on the active sites by physical (Physisorption) or chemical adsorption (Chemisorption). Physisorption is the interaction between the adsorbate and the active site through the weak van der Waals forces while in the chemical adsorption, the interaction results in the formation of a strong chemical bond. Physisorption is reversible and leads to the formation of multimolecular layers while chemisorptions are irreversible and form a unimolecular layer.

#### **3. Common adsorbents**

The applicability of the adsorbent in the removal of pollutants majorly depends on the adsorbent's characteristics. Some of the major characteristics include porosity, pore configuration, and the general nature of the surface of the adsorbent. The adsorption sites are spread throughout the solid. Adsorption sites'sizes are categorized into three: macropores, mesopores, and micropores. While macropores' diameter is larger than 50 nm, mesopores have a diameter that varies between 2 and 50 nm. Micropores' diameter is less than 2 nm. Many adsorbents are either naturally occurring or manufactured. Some of the common adsorbents include activated carbon, silica gel, zeolites, clay, nanoparticles, biological wastes, exchange resins, and water hyacinth.

#### **3.1 Activated carbon**

It is an organic material that forms a porous medium for adsorption. The structure of this medium is composite. The building blocks in this structure are carbon atoms.

*Adsorption and Its Applications: Using Zinc Adsorption on Water Hyacinth to Elaborate… DOI: http://dx.doi.org/10.5772/intechopen.104293*

Activated carbon can be used to remove substances such as dyes and pesticides as well as in the purification of wastewater [7]. This is because it is extremely effective in cadmium, lead, and zinc removal. The removal of pollutants such as chlorinated hydrocarbons, purification of helium, removal of phenols, removal of gas odors, and removal of nitrogen from the air have all utilized activated carbon [8].

#### **3.2 Silica gel, zeolite, clay minerals, and oxides**

Silica gel is a porous form of silica. It's mainly used to control moisture because of its high ability to absorb water vapor [9]. Also, it is used to control moisture in shoe boxes and to remove moisture from transformer oils and gases. Silica gel is widely used since it is cheap and can be regenerated [10]. Zeolite occurs naturally in the form of crystalline alumina that can be used in the adsorption of organic molecules from a gaseous phase. Zeolites have high water uptake since they have extensive surface area. Zeolite is used in industries to purify hydrogen gas and to recover carbon dioxide. In petroleum manufacturing, normal paraffin is separated from branched paraffin by the use of zeolites through adsorption. Clinoptilolite and bentonite, which are forms of zeolite, have been reported as proficient adsorbents in purification since they adsorb heavy metals [11, 12]. Clay minerals and oxides are available in nature and widely used because of their ability to adsorb many species of element. They adsorb cations such Cu2+, anions such as N2+, and neutral metallic species. Clay can be categorized into four: mica, smectite, kaolinite, and montmorillonite. Although they are readily available, they are less efficient in the adsorption of heavy metals when compared with zeolites.

#### **3.3 Nanomaterial, agricultural by-products, and biological waste**

Nanomaterial is a porous material whose pore diameter is less than 200 nm. These materials pose special properties such as distinctive surface and structural properties such as crystallinity and defect. They are used in processes such as ion exchange, catalysis, and separation. Nanomaterials are efficient in adsorption although they are expensive [13]. Most of the nanomaterials that have been used as an adsorbent include carbon nano-tube, activated carbon, and graphene. Agricultural by-products and biological wastes have been utilized in the adsorption of heavy metals. They are readily available, require modest processing, economical, offer selective adsorption and easy to generate. Agricultural by-products such as pecan shells, coconut shells, rice husks, cow dung, and maize cob have all been used effectively [14].

#### **3.4 Ion exchange resins**

These are organic materials (polymers) that can substitute ions within them with ions in a solution. This occurs when a solution containing ions is passed through the polymer. Resins can either be anionic or cationic. Anionic resins are negatively charged. As the solution passes through the resin, the positively charged ions in the solution are trapped since the resin is negatively charged [15]. Anionic resins are either weak or strong acids. Cationic resins are positively charged. They trap the negatively charged ions in the solution. The ion exchange resin is used in water softening to substitute Mg2+ and Ca2+ with Na+ converting hard water to soft water [16]. In this case, the resin is regenerated by rinsing it with a solution whose concentration of sodium ions is high. Resins are also used to purify water. In such a case, the poisonous heavy metal ions are replaced with ions such as sodium. Water containing no mineral content is purified using a resin that contains H+ and OH� to replace anions and cations.

#### **3.5 Water hyacinth and other low-cost adsorbents**

The plant can be grown in contaminated soils to absorb the heavy metal ions [17]. Also, it can be harvested, dried and ground into powder before dispensing it in heavy metal contaminated water. Water hyacinth powder should be allowed sufficient time in the contaminated water for adsorption to occur before the water can be used. The constituents of water hyacinth are cellulose (30–50%), hemicelluloses (20–40%), and lignin (15–30%) [18]. The cellulose contains functional groups such as O-H, which are involved in adsorption through deprotonation [19]. A comparative study on the efficiency of water hyacinth, water lettuce, and vetiver grass showed that the three plants have different abilities to eliminate water contaminants and their capacities are influenced by factors such as climate and temperature etc. [20]. Low-cost adsorbents provide cheap and readily available material to use as adsorbents. Lately, scientists have intensified research on the use of low-cost adsorbents in the removal of heavy metals from wastewater. Biological resources such as the agricultural waste have been put to the test and have proved useful. Rice husks, cow dung, sugarcane bagasse, sawdust, cashew nutshell, soybean hull, coconut shell, cotton hull, orange and banana peel have all successfully been used to remove heavy metals from wastewater. These adsorbents are cleaned and ground to the desired particle sizes before being used for adsorption while others are modified using modification techniques to improve the active sites [21].

#### **4. Adsorption kinetics**

During adsorption, equilibrium is not established immediately. The particle transfer from the solution to the adsorption site is restricted by the mass transfer resistance. The graph showing rate of adsorption with time is referred to as adsorption kinetics. It shows the rate of withholding or discharge of adsorbate from the aqueous solution to the adsorbent surface. The rate is influenced by the amount of the adsorbent, temperature, pH and particle size of the adsorbent, among others. The adsorption rate is restricted by resistance to diffusion in the solution (as adsorbate diffuses from the solution to the surface of the adsorbent) and within the adsorbent layers (as the adsorbate diffuses to the active site within the adsorbent) [22]. Kinetics studies are, therefore, important in determining the rate-limiting stage. Kinetics of adsorption is studied using kinetic models such as pseudo-first-order and pseudo-second-order.

#### **4.1 Pseudo-first-order**

To be able to analyze adsorption, several models have been applied. One such model is the pseudo-first-order. The linear form of the pseudo-first-order equation is [23]

$$\text{Log}\left(\mathbf{q}\_{\text{e}} - \mathbf{q}\_{\text{t}}\right) = \log\left(\mathbf{q}\_{\text{e}}\right) - \frac{\text{K1}}{2.303}t \tag{1}$$

*Adsorption and Its Applications: Using Zinc Adsorption on Water Hyacinth to Elaborate… DOI: http://dx.doi.org/10.5772/intechopen.104293*

Where qe (mg/g)—the amount of adsorbate on adsorbent at equilibrium, qt (mg/g)—the amount of heavy metal adsorbed on water hyacinth at a time t while K1(Min�1 ) is a rate constant of adsorption for pseudo-first-order. The values of qe and qt are determined using the adsorption capacity Eqs. (2) and (3) respectively [24].

$$\mathbf{p}\left(\mathbf{q}\_{\text{e}}\right) = \frac{\left(\mathbf{C1} - \mathbf{Ce}\right)}{M} \times \mathbf{V} \tag{2}$$

$$\mathbf{p}\left(\mathbf{q}\_{\text{t}}\right) = \frac{\left(\mathbf{C1} - \mathbf{C2}\right)}{M} \times \mathbf{V} \tag{3}$$

Where C1 and C2 refer to the initial and final concentration, Ce is the concentration of the analyte at equilibrium, M-mass of the adsorbent, V is the volume used. The adsorption rate constant (K1) for pseudo-first-order is computed from the gradient of a linear plot of log (qe-qt) versus t while qe is the value at the intercept.

#### **4.2 Pseudo-second-order**

The linear form of pseudo-second-order is [25];

$$\frac{\text{t}}{\text{qt}} = \frac{1}{\text{qe}}t + \frac{1}{\text{K2qe2}}\tag{4}$$

Where qe (mg/g)—the amount of adsorbate on adsorbent at equilibrium, qt (mg/g)—the amount of heavy metal adsorbed on water hyacinth at a time t while K2(Min�1 ) is a rate constant of adsorption for pseudo-second-order. The values of qt are determined using the adsorption capacity Eq. (2). The adsorption rate constant for pseudo-second-order (K2) and qe are computed from the gradient and intercept of a linear plot of t/qt versus t respectively. The linear plot with the highest R2 value is considered to describe the reaction best and is taken as the correct reaction order.

#### **4.3 Experimental example**

Adsorption studies of water hyacinth powder in the removal of Zn2+ (95.5 ppm) from an aqueous solution were conducted, and the data obtained analyzed using both pseudo-first and pseudo-second-order.

To prepare the stock solution, the following procedure was followed [26]:

• The amount of salt dissolved in a liter of distilled water to make 1000 ppm aqueous solutions of zinc was calculated using Eq. (5) [27].

$$\mathbf{m} = \frac{\mathbf{M} \mathbf{w}}{\mathbf{A} \mathbf{w}} \ge \frac{\mathbf{100}}{\mathbf{P}} \ge \frac{\mathbf{V}}{\mathbf{1000}} \tag{5}$$

Where m = mass (g) of analytical grade zinc nitrate, which was weighed, MW = molecular weight, Aw = Atomic mass, V = volume of the stock solution to be made, P = percentage purity of the salt.

$$\text{m} = \frac{297.48}{65.4} \text{ x} \frac{100}{95.97} \text{ x} \frac{1000}{1000} = 4.7396 \text{g}$$

$$\mathbf{C}\_1 \mathbf{V}\_1 = \mathbf{C}\_2 \mathbf{V}\_2 \tag{6}$$

Water hyacinth stems were collected from the shores of Lake Victoria, cut into smaller pieces, and cleaned vigorously with water to get rid of dust and other contaminants. To further prepare water hyacinth pieces for adsorption, the following procedure was followed

#### *4.3.1 Batch experiments*

To carry out the batch studies, the following procedure was followed [29]:

**Table 1** gives information about the initial and final analyte concentration (C0, and Ct), the volume of the solution (V), the mass of the adsorbent (M), the amount of adsorbate on water hyacinth at equilibrium (qe - mg/g) and at time t (qt-mg/g), the quotient t/qt, the difference between qe-qt and their logarithmic difference (log qe-qt).

Time and logarithmic differences presented between in **Table 1** were used to generate **Figure 1**. **Figure 1** shows the variation of the logarithmic difference (log qeqt) with time. The logarithmic difference (log qe-qt) in adsorption capacity at equilibrium and at particular time t decreased linearly as the adsorption time increased as shown between **Figure 1**.

Time and the quotient t/qt presented in **Table 1** were used to generate **Figure 2**. **Figure 2** demonstrates how the quotient t/qt changes with time t. The quotient of t /qt increased linearly with the adsorption time as shown in **Figure 2**.

*Adsorption and Its Applications: Using Zinc Adsorption on Water Hyacinth to Elaborate… DOI: http://dx.doi.org/10.5772/intechopen.104293*

#### **Table 1.**

*Kinetic data for Zn2+.*

The essential information presented in **Figures 1** and **2** was used in determining the kinetic parameters, which were used in establishing the most suitable model that could be used to describe the adsorption of Zn2+ on water hyacinth powder.

#### *4.3.2 Pseudo-first-order and pseudo-second-order kinetic parameters*

The pseudo-first-order and pseudo-second-order kinetic parameters include K1, qe, and K2. The initial concentration of Zn2+ was 95.5 ppm while the final

**Table 2.**

*The pseudo-first-order and pseudo second-order kinetic parameters of adsorption of zinc ions on powdered water hyacinth.*

concentration was 17.1 ppm. Using the Eqs. (1), (4) and the slopes in **Figures 1** and **2**, the pseudo-first-order and pseudo-second-order kinetic parameters can be calculated as shown:

a. For the pseudo-first-order

Comparing the slope from **Figure 1** with the linear form of pseudo-first-order equation (Eq. (1))

$$\mathbf{y} = -0.0209\mathbf{x} + 1.2253$$

$$\text{Log } (\mathbf{q}\_e - \mathbf{q}\_t) = -\frac{\mathbf{K1}}{2.303}t + \log\_e(\mathbf{q}\_e)$$

$$\mathbf{q}\_e = \mathbf{16.7996}, \text{while } \mathbf{k}\_1 = -2.303 \text{ x} - 0.0209 = 0.0481$$

#### b. For the pseudo-second-order

Comparing the equation from **Figure 2** with the linear form of pseudo-secondorder

$$\mathbf{y} = \mathbf{0}.\mathbf{0}\mathbf{5}\mathbf{7}\mathbf{8}\mathbf{x} + \mathbf{0}.\mathbf{7}\mathbf{0}\mathbf{3}\mathbf{3}$$

$$\frac{\mathbf{t}}{\mathbf{q}\mathbf{t}} = \frac{1}{\mathbf{q}\mathbf{e}}t + \frac{1}{\mathbf{K}\mathbf{2}\mathbf{q}\mathbf{e}\mathbf{2}}$$

qe = 1/0.0578, hence qe 17.3010, K2 = 0.0048. The pseudo-first-order and pseudo-second-order kinetic parameters are presented in **Table 2** which provides information about K1, qe, R2 , and K2.

The adsorption kinetics pointed out that the pseudo-first-order kinetic model expressed better the adsorption technique of zinc on water hyacinth with the regression coefficient R2 = 0.944 being higher than that of the pseudo-second-order (R2 = 0.8995).

#### **5. Adsorption isotherms**

Adsorption isotherm refers to a graph showing a relationship between adsorbate in the bulk and that on the surface of the adsorbent at a constant temperature. The adsorbate adsorption or release in an aqueous solution could be represented as follows [30]:

$$\text{S}\_{\text{f}}(\text{Me}) + \text{H}^{+} \text{ (solution)} \leftrightarrow \text{S}\_{\text{f}} \text{ (adsorbed)} + \text{Me} \text{(solution)}\tag{7}$$

*Adsorption and Its Applications: Using Zinc Adsorption on Water Hyacinth to Elaborate… DOI: http://dx.doi.org/10.5772/intechopen.104293*

where Sf represents the different adsorption sites on the adsorbent where the metal, Me can be retained. The adsorption of the metal on the sites could be enhanced by both the physical and chemical characteristics of the medium [31]. Langmuir and Freundlich's isotherms have been applied to heavy metal adsorption studies. The Langmuir isotherm assumes that adsorption is a reversible process and the adsorbing material has a definite number of active spots, which are evenly distributed [32]. The Freundlich isotherm describes multilayer adsorption [33]. The adsorption isotherms and modeling are very important in determining the accuracy of the adsorption process.

#### **5.1 Langmuir isotherms**

Langmuir adsorption was intended to illustrate gas–solid adsorption. However, it is also used to compute and compare the adsorption capacities of various adsorbents. Langmuir isotherm balances both adsorption and desorption rates. Adsorption is the measure of the portion of the adsorbent's surface that is open while desorption accounts for the portion of the surface of the adsorbent that is occupied [33]. The linear form of Langmuir isotherm that was used is shown in Eq. (8) [34]

$$\frac{1}{qe} = \left(\frac{1}{qmb}\right) \cdot \frac{1}{Ce} + \frac{1}{qm} \tag{8}$$

where: qe (mg/g)—equilibrium adsorption capacity, Ce (mg/l)—the amount of adsorbed adsorbate at equilibrium, qm (mg/g)—the highest amount of the adsorbate for every unit weight of adsorbent while b (l/mg)—Langmuir constant (binding affinity). The qm and b values are determined graphically from a plot of 1/qe against 1/Ce. The dimensionless factor RL, also known as the separation factor, is used to describe the Langmuir isotherm. It is determined using Eq. (9) [35]

$$\mathbf{R}\_{\rm L} = \frac{1}{b \ Co + 1} \tag{9}$$

where; RL < 1—adsorption is favorable, RL > 1—adsorption is unfavorable RL = 1—linear RL > 1—adsorption is irreversible

#### **5.2 Freundlich isotherms**

Freundlich isotherm equation is based on the fact that adsorption occurs on a heterogeneous surface. It is an experiential model, which takes into account the adsorptive active sites and their energy exponentially. The expression also considers the heterogeneity of the adsorbent's surface. It is expressed as shown in Eq. (10) [36]

$$\mathbf{q}\_{\mathbf{e}} = \mathbf{K}\_{\mathbf{F}} \mathbf{C}\_{\mathbf{e}}^{\mathbf{1}/\mathbf{n}} \tag{10}$$

The linear form of the Freundlich isotherm equation is shown in Eq. (11) [37].

$$\log \text{ q}\_{\text{e}} = \frac{\mathbf{1}}{n} \text{ log } \text{C}\_{\text{e}} + \log \text{ K}\_{\text{f}} \tag{11}$$

Where qe (mg/g)—the amount of the adsorbate adsorbed per unit weight of water hyacinth bio-material, Ce (mg/L)—the amount of unadsorbed adsorbent in the solution, Kf—a constant indicating adsorption capacity while n—adsorption intensity. Adsorption studies of zinc ions conducted using synthesized magnetite and baobab composite showed that the equilibrium data were suitably expressed by Freundlich because of the high correlation coefficient [38].

#### **5.3 Experimental example**

Equilibrium adsorption studies of water hyacinth powder in the removal of Zn2+ from aqueous solution were performed, and the data obtained were analyzed using both Langmuir and Freundlich isotherms. The working solutions of concentrations 0.5, 1.0, 10, 20, 40, 50, 60, 80, 100 ppm were prepared through serial dilution of the stock solution prepared in Section 4.3 following the formula indicated in Eq. (6).

#### *5.3.1 Batch experiments*

100 ml of aqueous solutions containing zinc metal ion (0.5–100 ppm) were added to 0.5 g of water hyacinth powder (<300 μm) in separate 500 ml beakers. The pH of the solutions was set at 5, and the solutions were stirred for 5 minutes. Adsorption was allowed to occur for 120 minutes. These experiments were performed at room temperature until the equilibrium was established. The mixture was filtered using gravity filtration (using Whatman Filter Paper no. 40 and plastic filter funnels). The residual concentrations of zinc ions in the filtrate were determined using the Atomic Absorption Spectrometer (AA-6300 SHIMADZU AAS). The data obtained from the adsorption studies were computed using Eqs. (2), (8) and (11), and tabulated in **Table 3**. **Table 3** gives information about the initial and equilibrium analyte concentration (C0, and Ce), reciprocal of equilibrium concentration (1/Ce), the amount of zinc ion on water hyacinth at equilibrium (qe - mg/g) and its reciprocal (1/qe), and the logarithmic values of both qe and Ce (log qe, log Ce).

The logarithmic values of Ce and qe accessible in **Table 3** were used to generate **Figure 3** which demonstrates demonstrates the variation of log qe with log Ce.

#### **Table 3.** *Equilibrium data for Zn2+.*

*Adsorption and Its Applications: Using Zinc Adsorption on Water Hyacinth to Elaborate… DOI: http://dx.doi.org/10.5772/intechopen.104293*

**Figure 3.** *Linearized Freundlich plot for the adsorption of Zinc ions onto water hyacinth powder.*

The logarithm of qe (adsorption capacity at equilibrium) increased linearly to the logarithm of Ce (concentration at equilibrium) as shown in **Figure 3**.

The quotients 1/Ce and 1/qe presented in **Table 3** were used to generate **Figure 4** which illustrates how the quotient 1/Ce varies with 1/qe. The reciprocal of qe (adsorption capacity at equilibrium) increased linearly with the reciprocal of Ce (concentration at equilibrium) as shown in **Figure 4**.

The information presented in **Figures 3** and **4** was used in computing the Freundlich and Langmuir parameters essential in determining the most suitable isotherm for explaining the adsorption of Zn2+ on water hyacinth powder.

#### *5.3.2 Freundlich and Langmuir parameters*

The Freundlich and Langmuir parameters include qm, b, n, and Kf. The initial concentration of Zn2+ was 100 ppm while the equilibrium concentration was 20.7 ppm. Using the Eqs. (8), (11) and the slopes from **Figures 3** and **4**, the Langmuir and Freundlich parameters were calculated as shown;

#### a. Freundlich parameters

Comparing Eq. (11) with the equation from **Figure 3**

**Figure 4.** *Linearized Langmuir plot for the adsorption of Zinc ions onto water hyacinth powder.*

#### **Table 4.**

*Langmuir and Freundlich isotherms parameters for adsorption of zinc ions from aqueous solution ground using water hyacinth powder.*

> *log* Kf ¼ 0*:*3276 hence Kf ¼ 0*:*3487 1*=*n ¼ 1*:*4944 hence n ¼ 0*:*6692

b. Langmuir parameters

Comparing Eq. (8) with equation from **Figure 4**;

$$\mathbf{y} = 0.5636\mathbf{x} + 0.0452$$

$$\frac{\mathbf{1}}{\mathbf{q}\mathbf{e}} = \left(\frac{\mathbf{1}}{\mathbf{q}\mathbf{m}\mathbf{b}}\right) \cdot \frac{\mathbf{1}}{\mathbf{C}\mathbf{e}} + \frac{\mathbf{1}}{\mathbf{q}\mathbf{m}}$$

$$\frac{\mathbf{1}}{\mathbf{q}\mathbf{m}} = 0.0452 \text{ hence } \mathbf{q}\_{\mathbf{m}} = 22.1239$$

$$\frac{\mathbf{1}}{\mathbf{q}\mathbf{m}.\mathbf{b}} = 0.5636 \text{ replacing } \mathbf{q}\_{\mathbf{m}} \text{ we get } \mathbf{b} = 0.0802$$

From Eq. (9), RL= 1 *b Co*þ1 *:* The value of RL = 1 ð0*:*0802x 100Þþ1 = 0.1109.

The Freundlich and Langmuir adsorption parameters are presented in **Table 4** which provides information about n, b, Kf, qm, and the regression coefficient of **Figures 3** and **4**.

The use of water hyacinth powder in the adsorption of zinc ions correlated well with the Langmuir model in contrast to the Freundlich model since it had the highest regression coefficient (R2 = 0.975). The linearity of **Figure 4** demonstrates the validity of the Langmuir isotherm whose basis is the formation of a monolayer on the surface of water hyacinth powder. Results similar to these were obtained by [39] where the equilibrium data during adsorption of Zinc (II) ions from aqueous solution using functionalized lignocelluloses derived from waste biomass was suitably represented by Langmuir isotherm due to high correlation coefficients. The adsorption process was favorable since the separation factor RL (0.1109) is less than 1. The Freundlich parameter *n* (n = 0.6692) indicates that the adsorption (physical adsorption) of zinc is unfavorable since 2 < *n* < 10 indicates favorable adsorption. The Langmuir constant b is important in establishing the adsorption affinity that relates the bond energy between adsorbent and adsorbate [40].

#### **6. Thermodynamics of adsorption**

To evaluate the adsorption process further, it is imperative to examine the adsorbate's binding thermodynamics on the adsorbent. Thermodynamics include changes

*Adsorption and Its Applications: Using Zinc Adsorption on Water Hyacinth to Elaborate… DOI: http://dx.doi.org/10.5772/intechopen.104293*

in entropyð Þ Δ*S*° , enthalpy ð Þ Δ*H*° , and Gibbs free energy (Δ*G*°Þ, which are computed using Eqs. (12)–(14).

The Gibbs free energyΔG°is given by [41];

$$
\Delta \mathbf{G}^{\circ} = -\mathbf{R} \, \mathbf{T} \, \ln \, \mathbf{K} \mathbf{c} \tag{12}
$$

Where ΔG° (kJ/mol) is the change in the Gibbs free energy of adsorption, R is the gas constant (8.314 J/ (Kmol)), T is the temperature in Kelvin, Kc is the equilibrium constant. Equilibrium constant Kc can be stated with reference to Gibbs free energy ΔG°and entropy ΔS° as indicated in the Van't Hoffs reaction given in Eq. (13) [42].

$$
\ln \text{ Kc} = -\frac{\Delta \text{H}^{\circ}}{\text{R}} \left( \frac{\text{1}}{\text{T}} \right) + \frac{\Delta \text{S}^{\circ}}{\text{R}} \tag{13}
$$

Where is ΔH° (kJ/mol) enthalpy change while ΔS° (J/ (mol K) is entropy change. Kc is computed by following Eq. (14) [43].

$$\mathbf{Kc} = \frac{\mathbf{Qe}}{\mathbf{Ce}} \tag{14}$$

Where Qe is the equilibrium adsorption capacity while Ce is the equilibrium concentration of the heavy metal under investigation.

To determine thermodynamic parameters of adsorption, the following procedure is followed;

**Table 5.**

*Gibbs free energy (ΔG°), enthalpy (ΔH°), entropy (ΔS°) and spontaneity of an adsorption process.*

Negative values of ΔG°indicate that the process is spontaneous and will support the forward reaction while positive values show the process is nonspontaneous, hence will support the reverse reaction [44]. The spontaneity of the adsorption process depends on the values of ΔH°andΔS° as outlined in **Table 5**. **Table 5** outlines information about Gibbs free energy (ΔG°) and how the spontaneity of an adsorption process is influenced by both enthalpy (ΔH° and entropy (ΔS°).

When ΔH° >0, then the adsorption process is endothermic and requires to absorb heat from the environment. When ΔH° < 0, the adsorption process is exothermic and releases heat into the surroundings [45]. If ΔS° < 0, it indicates that the adsorption is orderly but if ΔS° > 0 shows adsorption is disorderly at the surface of the adsorbent [46].

#### **7. Application of adsorption**

#### **8. Conclusion**

In the review, we have examined the adsorption process, adsorbent, adsorption mechanism, and some of the common adsorption applications. Also, we have scrutinized essential kinetic and thermodynamic parameters of adsorption. Kinetic models and isotherms described include pseudo-first-order, pseudo-second-order, Langmuir, and Freundlich. Kinetic parameters explained include adsorption capacity, pseudofirst and pseudo-second constants while the isotherm parameters described herein include adsorption intensity (n), binding affinity (b), Freundlich constant (Kf),

*Adsorption and Its Applications: Using Zinc Adsorption on Water Hyacinth to Elaborate… DOI: http://dx.doi.org/10.5772/intechopen.104293*

adsorption capacity (qm), separation factor (RL). We have used the adsorption of zinc ions to illustrate how kinetic and isotherm parameters are computed during adsorption studies. The adsorption data for zinc from the aqueous solution using water hyacinth fitted well in the Langmuir model (R2 = 0.975) in comparison to the Freundlich model (R2 = 0.9659) based on the correlation coefficients. Additionally, the study showed that zinc adsorption obeys the pseudo-first-order reaction kinetic model. The adsorption of zinc ions using water hyacinth powder was favorable since the separation factor RL was less than 1. The Freundlich parameter *n* an indicator of the probability of physical adsorption taking place showed that physisorption was unfavorable since *n* < 2. Further, this chapter clearly describes how thermodynamics parameters (entropyð Þ Δ*S*° , enthalpy ð Þ Δ*H*° , and Gibbs free energy (Δ*G*°Þ) are evaluated during adsorption studies. Also, the chapter details the meaning of entropyð Þ Δ*S*° , enthalpy ð Þ Δ*H*° , and Gibbs free energy (Δ*G*°Þ concerning the spontaneity of the adsorption process.

#### **Acknowledgements**

I am deeply obliged to my supervisors, Dr. John Onam Onyatta and Prof. Amir O. Yusuf, for the invaluable comments, suggestions, corrections, patience, and constructive criticisms that went a long way to refine this work. Without whose help, this work would not have been a success. Thank you very much for making sure I kept my focus.

#### **Originality/declaration**

The author declares that this work is their original work and has not been submitted elsewhere for examination or publication.

#### **Conflict of interest**

The author declares no conflict of interest.

#### **Author details**

Joseph Munene Mwaniki University of Nairobi, Nairobi, Kenya

\*Address all correspondence to<EMAIL_ADDRESS>
© 2022 The Author(s). Licensee IntechOpen. This chapter is distributed under the terms of the Creative Commons Attribution License (http://creativecommons.org/licenses/by/3.0), which permits unrestricted use, distribution, and reproduction in any medium, provided the original work is properly cited.

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[35] Akrawi HSY, Al-Obaidi MA, Abdulrahman CH. Evaluation of Langmuir and Frendlich isotherm equation for zinc adsorption in some calcareous soil of Erbil province north of Iraq. IOP Conference Series: Earth and Environmental Science. 2021;**761**(1):012017. DOI: 10.1088/ 1755-1315/761/1/012017

[36] Nnaji CC, Agim AE, Mama CN, Emenike PC, Ogarekpe NM. Equilibrium and thermodynamic investigation of biosorption of nickel from water by activated carbon made from palm kernel chaff. Scientific Reports. 2021;**11**(1): 7808. DOI: 10.1038/s41598-021-86932-6

[37] Palanivell P, Ahmed OH, Latifah O, Abdul Majid NM. Adsorption and desorption of nitrogen, phosphorus, potassium, and soil buffering capacity following application of chicken litter biochar to an acid soil. Applied Sciences. 2019;**10**(1):295. DOI: 10.3390/ app10010295

[38] Abdus-Salam N, Adekola SK. Adsorption studies of zinc(II) on magnetite, baobab (Adansonia digitata) and magnetite–baobab composite. Applied Water Science. 2018;**8**(8):222. DOI: 10.1007/s13201-018-0867-7

*Adsorption and Its Applications: Using Zinc Adsorption on Water Hyacinth to Elaborate… DOI: http://dx.doi.org/10.5772/intechopen.104293*

[39] Wadhawan S, Jain A, Nayyar J, Mehta SK. Role of nanomaterials as adsorbents in heavy metal ion removal from waste water: A review. Journal of Water Process Engineering. 2020;**33**: 101038. DOI: 10.1016/j. jwpe.2019.101038

[40] Togue Kamga F. Modeling adsorption mechanism of paraquat onto Ayous (Triplochiton scleroxylon) wood sawdust. Applied Water Science. 2019; **9**(1):1. DOI: 10.1007/s13201-018-0879-3

[41] Dang J, Wang H, Wang C. Adsorption of toxic zinc by functionalized lignocellulose derived from waste biomass: Kinetics, isotherms and thermodynamics. Sustainability. 2021;**13**(19):10673. DOI: 10.3390/ su131910673

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#### **Chapter 6**

## Adsorption of Chromium from an Aqueous Solution onto Chitosan Beads Modified with Sodium Dodecyl Sulfate (SDS)

*Naoki Kano, Zou Ming, David Eva Vanessa Anak and Muhammad Nabil Md Sari*

#### **Abstract**

The goal of this research is to make chitosan beads that have been treated with sodium dodecyl sulfate (SDS) to remove chromium (Cr) from an aqueous solution effectively. The successful synthesis of the SDS-chitosan was proven through characterization, which were carried out using by scanning electron microscopy–energy dispersive X–ray spectroscopy (SEM-EDS), Fourier transform-infrared spectroscopy (FT-IR) and X-ray photoelectron spectroscopy (XPS). The adsorption of Cr on the SDS material was investigated by varying experimental conditions such as pH, contact time and adsorbent dosage. The maximum adsorption capacity of SDS-chitosan for Cr (III) was estimated to be 3.42 mgg1 and 3.23 mgg1 for Cr(VI). Based on the results of adsorption kinetics and isothermal models, the adsorption process conform to the pseudo-second-order and Langmuir isotherm models. This indicates that the adsorption of Cr on SDS-chitosan is mainly dominated by chemical adsorption and monolayer reaction. In addition, according to thermodynamic analyses, the adsorption of Cr is an endothermic reaction. These results show that the new adsorbent has obvious application prospect for removing Cr.

**Keywords:** chitosan beads, sodium dodecyl sulfate (SDS), chromium, adsorption, adsorption isotherms, adsorption kinetics

#### **1. Introduction**

Due to the rapidly growing number of manufacturing industries, toxic metal contamination in aquatic environments has gotten a lot of attention. Among the contaminants, heavy metals are targeted for major environmental concern because they are non-biodegradable, and they cannot be decomposed or metabolized [1]. Several metals cause serious health and environmental problems, and chromium (Cr) compounds are one of the most toxic contaminants in wastewater due to their high solubility and toxicity, as well as their free transferability [2].

Cr has been widely applied in a variety of industrial activities due to its excellent properties, including electroplating, leather tanning, nuclear power plants, and textile industries [3, 4]. Furthermore, it can be used for anodizing, corrosion control, and chemical manufacturing [5–7]. In a natural environment, Cr usually exists in two stable oxidation states: trivalent Cr(III) and hexavalent Cr(VI). Meanwhile, other oxidation states are not stable in aerated aqueous media [8]. Specifically for Cr(VI), it may exist in the form of CrO4 2 or HCrO4 in a natural aqueous environment, whereas Cr(III) is inclined to form [Cr(H2O)6] 3+, Cr(H2O)5(OH)2+, Cr(H2O)4(OH)2+, or Cr(III) organic complexes.

Given its considerable risk to biological systems, many studies have focused on the removal of Cr(VI). Cr(VI) is highly toxic, carcinogenic, and mutagenic [8, 9].  On the other hand, Cr(III) is the most stable form in reducing conditions and it exists as cationic species Cr(OH)2 + and Cr(OH)2 + , with the first or second hydrolysis products dominating at pH values ranging from 4 to 8. Although Cr(III) is an essential microelement for the effective maintenance of mammal's glucose, lipid, and protein metabolism [10], high doses of Cr(III) may cause negative consequences to the environment. Moreover, there are currently just a few articles on the adsorption of Cr(III). Therefore, the development of a recovery method for this metal (both Cr(III) and Cr(VI)) is significant from an environmental aspect.

Ion exchange, precipitation, ultrafiltration, reverse osmosis, and electro dialysis are one of the physical and chemical technologies that have been reported for the removal of heavy metals [11]. However, these procedures have some drawbacks, such as a high consumption of reagents and energy, low selectivity, high operational costs, and difficult further treatment due to toxic sludge production [12]. Adsorption is an effective method for removing metallic ions from aqueous solutions [10, 13]; and biological adsorption (biosorption) is one of the most environmentally friendly, costeffective, recyclable, and technically simple methods [14, 15].

Among the many biosorbents available, chitosan can be an excellent biosorbent for metals because its amine (-NH2) and hydroxyl (-OH) groups may serve as coordination sites to form complexes with various heavy metal ions [16]. Chitosan, whose full chemical name is (1,4)-2-amino-2- deoxy-β-D-glucose, has been proven to be particularly effective as a biosorbent for the recovery of several toxic metals, including mercury (Hg), uranium (U), molybdenum (Mo), vanadium (V), and platinum (Pt) [17–19]. It can be employed as an environmentally friendly adsorbent because it is cost effective, and it does not result in secondary pollution. Chitosan is a polymer that is made via alkaline deacetylation of chitin, which comes from cellulose, the most abundant biopolymer. It can be acquired from the shells of seafood, such as prawns, crabs, and lobsters [20]. The biopolymer has a high nitrogen content, which is present in the form of amine groups, free amino groups, and hydroxyl groups, all of which are responsible for metal ion binding through chelation mechanisms [21].

Despite the uses of chitosan in the removal of various pollutants have been adequately reviewed [22], on the other hand, it has some defects, including notable swelling in aqueous media and nonporous structures, resulting in a very low surface area [23]. Therefore, a variety of chemical modifications can be used to produce chitosan derivatives that improve the removal efficiency of the heavy metal [24]. For example, to increase the number of exposed active sites, several chemical or physical modifications can be adopted [25, 26]. Moreover, since silicon dioxide has numerous properties, including rigid configuration, porosity, and large surface area, it can be

*Adsorption of Chromium from an Aqueous Solution onto Chitosan Beads Modified with Sodium… DOI: http://dx.doi.org/10.5772/intechopen.104093*

employed to counteract the defects of chitosan. In addition, modified silicon dioxide has been produced through a graft between silanol groups and ligands [27–29]. In prior work, we synthesized a hybrid membrane of carboxymethyl chitosan and silicon dioxide as adsorbents for the removal of Cr(VI) [30]. Furthermore, we used epichlorohydrin (EP) and glutaraldehyde (GA) as cross-linked agents in an adsorption experiment of chromate ions onto cross- linked chitosan [31].

In this work, we evaluated the adsorption of chitosan modified with sodium dodecyl sulfate (SDS) as part of the adsorption study of Cr using modified chitosan. SDS-modified chitosan beads have been reported to be effective for removing cationic dyes [32]. Adsorption on surfaces is enabled by the metal ion strength and the presence of key functional groups on the polymer chain [33–35]. The particle aggregation via a bridging structure can be described as a two-step pathway: (1) initial chain adsorption and bridging, followed by (2) floc maturation/ reconfiguration. Before the interparticle connection occurs, the chain of SDS must be adsorbed on a chitosan surface [36]. Furthermore, chitosan modified with SDS has recently been used for the removal of heavy metals, such as cadmium [37, 38]. However, the use of SDSmodified chitosan as a Cr adsorbent, using different initial concentrations of SDS to optimize the adsorbent, has rarely been investigated. The objective of the present research is to determine the efficacy of SDS-modified chitosan beads as a sorbent for Cr(III) and Cr(VI) for future practical applications, as well as to understand the adsorption mechanism. After the characterization of SDS-chitosan by scanning electron microscopy-energy dispersive X-ray spectroscopy (SEM-EDS), X-ray photoelectron spectroscopy (XPS) and Fourier transform-infrared spectroscopy (FT-IR), batch experiments with SDS-modified chitosan beads were carried out to optimize the parameters and to obtain the maximum removal of Cr(III) and Cr(VI). In this study, we examined the effects of several parameters, including solution pH, contact time, adsorbent dosage, and initial concentration.

#### **2. Materials and methods**

Chemical reagents, such as chitosan and sodium dodecyl sulfate (SDS; M.W.: 288.372 g/mol), were purchased from Tokyo Chemical Industry Co., Inc. Acetic acid, NaOH, HNO3, NaSO4, ethylenediaminetetraacetic acid disodium salt dihydrate, and toluene, were purchased from Kanto Chemical Industry Co., Inc. All reagents used were of analytical grade. During the whole working process in this study, the water (>18.2MΩ) treated by the ultrapure water system (RFU 424TA, Advantech Aquarius) was employed. K2CrO7 standard solution (1000 mgL1 from Kanto Chemical Co., Inc.) was diluted and used to prepare the Cr standard solution for calibration.

#### **2.1 Synthesis of the adsorbent**

In this study, chitosan powder with a molecular weight (50–190 kDa) and degree of deacetylation (> 80%) was used. After drying, the viscosity of chitosan was 20 to 100 mPas (in 0.5%Acetic acid soln., 20°C). Firstly, 1.5 g of chitosan was placed in acetic acid solution (2.0%), and the solution was mixed for 24 h. The chitosan-gel was prepared by dropping the above chitosan solution into 200 mL of 0.20 molL1 NaOH. Consequently, the obtained gel was rinsed with ultrapure water until its pH reached 7 after stirring for 24 h. Secondly, 200 tablets of chitosan-gel beads were placed in 100 mL of SDS solution (including the fixed concentration of SDS), and then left for

**Figure 1.**

*Procedure employed for the synthesis of sodium dodecyl sulfate (SDS)-chitosan beads.*

five days. Thus, chitosan-gel beads modified with SDS were obtained, and finally, they were dried at 60°C overnight for use as an adsorbent. The synthesis procedure for the SDS-chitosan beads is demonstrated in **Figure 1**. It is considered that the prepared adsorbent has a bilayer of SDS over the surface of pure chitosan beads. This bilayer can have a higher ion capturing capacity [37].

#### **2.2 Characterization of the adsorbent**

The diameter of chitosan beads was measured to be about 0.5–2 mm (judging from 200 beads as representative chitosan beads). After weighing 200 hydrogel beads, the dry weight per chitosan bead was estimated to be 3.7 � 10�4 g, which suggests that the adsorbent contained 98% moisture. In order to determine the physicochemical properties of pristine and modified chitosan, several characterization methods have been employed. FT-IR spectra of the samples were recorded in the range of 4000–500 cm�1 with a JASCO Japan FTIR-4200 spectrophotometer using KBr pellet pressing method. The surface morphology and element distribution of the chitosan beads before and after the adsorption of Cr were observed using SEM-EDS (JEOL Japan: JCM-6000 with JED-2300). X-ray Photoelectron Spectroscopy (XPS, Thermo Scientific Center: K-Alpha) was also used to assess the surface chemistry properties of SDS-modified chitosan beads.

#### **2.3 Adsorption experiments**

The beads were put into contact with 50 mL of an aqueous solution containing Cr (III) or Cr (VI) ion with a known initial concentration. The pH of each solution was adjusted using 0.1 mol*=*L�1 NaOH and 0.1 mol*=*L�1 HNO3. The flask was then placed into an automatic shaker. In the adsorption experiments, the parameters were varied such as the pH range of 1�7, contact time of 1�72 h, SDS initial concentration of 10�9000 mg∙L�1 , adsorbent dosage of 0.01�0.06 g∙L�3 , temperature of 288�318 K, and initial Cr(III) or Cr(VI) concentration of 0.1�3.0 mg∙L�1 . The concentrations of Cr in the filtrate were measured by ICP-MS (Thermo Scientific Center: X-series II). The adsorption capacities of the chitosan beads modified with SDS were calculated by the following Eq. [9]:

$$q\_{\epsilon} = (\mathbf{C}\_{i} - \mathbf{C}\_{\epsilon}) \times \mathbf{V}/m \tag{1}$$

*Adsorption of Chromium from an Aqueous Solution onto Chitosan Beads Modified with Sodium… DOI: http://dx.doi.org/10.5772/intechopen.104093*

where *qe* represents the adsorption capacity at the equilibrium (mg∙g1 ), *Ci* and *Ce* are the initial and equilibrium concentrations of Cr in the batch system, respectively (mg∙L1 ), *V* is the volume of the solution (L), and *m* is the weight of the adsorbents (g).

#### **3. Result and discussion**

#### **3.1 Characterization of materials**

#### *3.1.1 SEM-EDS micrographs*

**Figure 2** shows the SEM-EDS images of the chitosan beads and SDS-chitosan. These images showed the unevenness of the surface. This was most likely due to the release of water from the adsorbent during the drying process of chitosan. In addition to the irregularities, SDS-chitosan showed a mesh-like pattern. This was thought to be due to SDS modified on the chitosan surface. In particular, the modified chitosan beads with a high SDS concentration stood out compared to the other beads. As the irregularities on the adsorbent's surface were considered, the adsorption proceeded in

#### **Figure 2.**

*Scanning electron microscopy (SEM) images of chitosan beads (a) and SDS-chitosan (b), and mapping images of SDS-chitosan after the adsorption of Cr(III).*

two ways: physical and chemical adsorption. From the mapping images, it was proven that Cr ions were adsorbed onto the adsorbent surface.

#### *3.1.2 FT-IR and XPS spectra*

The surface functional groups and the chemical compositions of the modified SDSchitosan beads were identified by FTIR and XPS analysis, respectively. The FT-IR results of chitosan and cross-linked chitosan are shown in our previous work [39]. **Figure 3** shows the FT-IR results of SDS-chitosan and the chitosan beads in this study, with the peaks of SDS-chitosan and the chitosan beads can clearly visible. It is apparent from this figure that the main peaks common to each adsorbent were due to the - OH group at 3400–3500 cm1 and the aliphatic methylene group at 2871 cm1 . The amine and ether groups are shown by wide peaks at 1560–1640 cm1 and 1110 cm1 , respectively. For SDS-chitosan, the peak at 1248 cm1 was characteristic of the asymmetrical vibration of the C-O-S group, confirming that the prepared adsorbent was a composite of SDS and chitosan.

The chitosan beads with varying initial loading concentrations of SDS were analyzed using XPS, as shown in **Figure 4**. The C1s spectra of these samples displayed peaks at 284.5, 286.5, and 288.5 eV, corresponding to C-C, C-O, and C=O bonds, respectively. The S2p spectra of SDS600 and 6000-chitosan displayed peaks at 169 eV. The S that seemed to be derived from SDS was also not detected at SDS concentrations of 0 and 100 mg/L but was detected at concentrations of 600 and 6000 mg/L.

#### **3.2 Adsorption experiment**

#### *3.2.1 Effect of the initial SDS concentration*

To determine the optimum initial loading concentration of SDS for Cr(III) removal, the chitosan beads were modified with SDS solutions ranging from 10 to 9000 mg/L. The adsorption experiments were performed under the following conditions: an initial concentration of Cr(III) of 1 mg/L, a contact time of 2 days, and adsorbent dosage of 0.05 g, a pH of 4, and a temperature of 25°C. In the meantime, initial SDS concentrations were varied from 10 to 6000 mg/L for Cr(VI) removal. The

**Figure 3.** *Fourier transform-infrared spectroscopy (FT-IR) spectra of chitosan and SDS-chitosan.*

*Adsorption of Chromium from an Aqueous Solution onto Chitosan Beads Modified with Sodium… DOI: http://dx.doi.org/10.5772/intechopen.104093*

**Figure 4.**

*X-ray photoelectron spectroscopy (XPS) spectra of chitosan (a), SDS100-chitosan (b), SDS600-chitosan (c), and SDS6000-chitosan (d) beads.*

experiment was performed under the following conditions: adsorbent dose of 0.05 g, initial concentration of Cr(VI) of 1 mg/L, pH of 4, and contact time of three days.

The results of Cr (III) adsorption are shown in **Figure 5**. The adsorption capacity continuously increased within 6000 mg/L as the SDS concentration increases, but the adsorption amount was almost constant at further higher concentrations. Thus, 6000 mg/L was the optimum initial loading SDS concentration for Cr(III). On the other hand, the adsorption capacity of SDS for the removal of Cr(VI) increased with an increase of the SDS concentration from 10 to 40 mg/L, and after that it decreased as shown in **Figure 6**. The maximum capacity for Cr(VI) was obtained at the initial SDS concentration of 40 mg/L.

**Figure 5.** *Effect of the initial SDS concentrations on the adsorption of Cr(III).*

#### **Figure 6.**

*Adsorption of Cr(VI) for different initial SDS concentrations (10–6000 mg/L). (a) Effect of initial SDS concentrations (10–6000 mg/L) on the adsorption of chromium (Cr) (VI). (b) Effect of initial SDS concentrations (10–1000 mg/L) on the adsorption of Cr(VI).*

#### *3.2.2 Effect of pH*

The effect of pH on Cr(III) adsorption by SDS-chitosan was investigated in the pH range of 4–7. Other parameters were set as the following: the contact time was 24 h, the temperature was 25°C, the adsorbent dosage was 0.4 mg/L, and the initial Cr(III) concentration was 1 mg/L. The results are shown in **Figure 7**. As shown in **Figure 7**, the adsorption capacity of Cr(III) increased with the increase of pH from 4 to 7. However, it was found that Cr (III) precipitated as Cr (OH)3 at the pH value in 6 and 7. Thus, a pH of 6 and 7 were not suitable for the adsorption experiments and optimized pH value was 4.

As for the adsorption experiment of Cr(VI), pH was set in the range of 4 to 10 (**Figure 8**). In this experiment, the shaking time was 24 h, the temperature was 25°C, and the dose of adsorbent was 0.02 g/L. As shown in **Figure 8**, pH affects Cr(VI)

**Figure 7.** *Effect of pH on the adsorption of Cr(III) by SDS-chitosan.*

*Adsorption of Chromium from an Aqueous Solution onto Chitosan Beads Modified with Sodium… DOI: http://dx.doi.org/10.5772/intechopen.104093*

**Figure 8.** *Effect of pH on the adsorption of Cr(VI) onto the SDS-modified chitosan beads.*

adsorption on SDS-modified chitosan beads. The acidity of the solution brought a significant effect on the adsorption of Cr(VI) on SDS-modified chitosan beads where the amino groups of the chitosan were protonated. Cr(VI) ions were most effectively adsorbed at pH 4–5, which may be related to changes in surface charge on the adsorbent. As pH increased above pH 5, the uptake decreased.

#### *3.2.3 Effect of contact time*

The influence of contact time on the adsorption of Cr(III) by SDS-chitosan was investigated. The experiment was conducted under the following conditions: a pH of 4, a temperature of 25°C, an adsorbent dosage of 0.05 g, and an initial concentration of Cr(III) of 1 mg/L. **Figure 9** shows that the adsorption capacity of SDS-chitosan for Cr (III) increased sharply within the first 24 h, and continued until the contact time reached 48 h. Thus, 48 h was selected as the optimized contact time.

We have also studied how the contact time affects the adsorption capacities of SDS-modified chitosan beads towards Cr(VI) with varying contact times from 1 to 96 h (**Figure 10**). In this experiment, the concentration of Cr(VI) was set as 1 mg/L with the dose of 0.05 g at a temperature of 25°C adjusted pH to 4.

**Figure 9.** *Effect of contact time on the adsorption of Cr(III) by SDS-chitosan.*

**Figure 10.** *Effect of the contact time on the adsorption of Cr(VI) on the SDS-modified chitosan beads.*

The adsorption capacity of SDS-chitosan beads for Cr(VI) increased sharply within the first 24 h, which may be attributable to the availability of the sites on the surface of the adsorbent. It is suggested that a concentration gradient is present in both the adsorbent and adsorbate in the solution [40]. Then, adsorption reached equilibrium at 72 h, and afterwards, there was no appreciable increase (**Figure 10**). Hence for further studies, the optimized contact time was taken as 72 h.

#### *3.2.4 Effect of the adsorbent dosage*

The adsorbent dosage is an important factor that affects the adsorption capacity. To study the effect of adsorbent dosage on the adsorption of Cr(III), adsorption experiments were conducted at a pH of 4, a temperature of 25°C, a contact time of 24 h, and an initial concentration of Cr(III) of 1 mg/L. The results are shown in **Figure 11**. The adsorption rate increased as the adsorbent dosage increased, reaching approximately 70% at 0.8 mg/L, after that, there was no appreciable increase. Thus, 1 mg/L was selected as the optimized adsorbent dosage.

Meanwhile for the adsorption experiment of Cr (VI), 0.8 mg/L was regarded as the optimum dosage. The experiments were performed by varying the dosage (from 0.4 to 1.0 mg/L) and keeping all other parameters constant (temperature: 25°C; pH: 4; contact time: 24 h; initial concentration: 1.0 mg/L). The results are shown in **Figure 12**.

**Figure 11.** *Effect of the adsorbent dosage on the adsorption of Cr(III) by SDS-chitosan.*

*Adsorption of Chromium from an Aqueous Solution onto Chitosan Beads Modified with Sodium… DOI: http://dx.doi.org/10.5772/intechopen.104093*

**Figure 12.**

*Effect of the adsorbent dosage on the adsorption of Cr(VI) on the SDS-modified chitosan beads.*

#### *3.2.5 Effect of competitive ions*

In this study, adsorption experiments of Cr(III) were conducted in the presence of several competitive ions with different concentrations (0, 50, 100, and 200 mg/L) of Na+ , Ca2+, Ba2+, K+ , Mg2+, and Sr2+. The initial concentration of Cr(III) was set to 1.0 mg/L, the pH was 4, the temperature was 25°C, the contact time was 24 h, and the adsorbent dosage was 1.0 mg/L. The effect of competitive ions on the adsorption of Cr (III) is shown in **Figure 13**. It was confirmed that the adsorption capacity of Cr(III) did not decrease at all even in the presence of other metal ions.

Conversely, it was found that Cr(VI) adsorption was inhibited when the concentrations of coexisting ions were high. The effect of competitive anions on the adsorption of Cr(VI) is shown in **Figure 14**. In this study, the initial concentration of Cr(VI) was fixed to 1 mgL1 . These counter ions were tested collectively, and all the ions were included at 50, 100, or 200 mg/L in solution. From this figure, the removal of Cr (VI) was remarkably reduced under the presence of common ions at above 50 mgL1 (i.e., 50 times the Cr(VI) concentration or more), although no substantial decrease was observed when the concentration of each common ion was less than 10 mgL1 in our previous preliminary experiments.

**Figure 13.** *The effect of competitive ions on the adsorption of Cr(III).*

**Figure 14.**

*Effect of competitive anions on the adsorption of Cr(VI) on the SDS-modified chitosan beads.*

#### **3.3 Adsorption isotherms**

Adsorption isotherms describe the interactive process between the adsorbents and adsorbates in aqueous medium at the attained saturation point. Adsorption isotherms of Cr(III) and Cr(VI) on SDS-chitosan were identified with different initial concentrations from 0.01 to 2 mg/L under optimized conditions for the pH (4), contact time (48 h for Cr(III), 72 h for Cr(VI)), and dosage of the adsorbent (1 mg/L for Cr(III), 0.8 mg/L for Cr(VI)). The adsorption of Cr on SDS-chitosan was evaluated using typical adsorption isotherms, the Langmuir and Freundlich models (**Figure 15(a)** and **(b)**).

The adsorption data obtained for Cr(III) using SDS-chitosan were analyzed by Langmuir and Freundlich equations (**Figure 16(a)** and **(b)**).

Langmuir equation:

$$\frac{C\_{\epsilon}}{q\_{\epsilon}} = \frac{C\_{\epsilon}}{q\_{\max}} + \frac{1}{K\_{L}q\_{\max}}\tag{2}$$

Freundlich equation:

$$
\log q\_{\varepsilon} = \lg \mathcal{K}\_{\mathcal{F}} + (\mathbf{1}/n)\lg \mathcal{C}\_{\varepsilon} \tag{3}
$$

where *Ce* and *qe* are the concentration of Cr at the equilibrium (mg∙L�1 ) and the amount of adsorption of Cr(III) at the equilibrium (mg∙g�1 ), respectively, *qmax* is the maximum adsorption capacity on the surface of the chitosan bead (mg∙g�1 ), *KL* is the Langmuir constant related to the adsorption strength or intensity (L∙mg�1 ), *KF* is the Freundlich constant and 1*=n* indicates the adsorption intensity of the system [41, 42].

The Langmuir isotherm model assumes that a monolayer adsorption occurs on the surface of an adsorbent. The slope of the linearized Langmuir isotherm can be used to explain the type of sorption using the Hall separation factor (*RL*), which is favorable (0 <*RL* <1), unfavorable (*RL* < *0*), linear (*RL* ¼ 1) or irreversible (*RL* ¼ 0) [40]. The Freundlich isotherm model assumes a multilayer adsorption process without considering the adsorption saturation, which occurs on the multilayer heterogeneous surface. The value of the adsorption intensity of the system 1ð Þ *=n* can be used to judge the difficulty of the adsorption process: irreversible 1ð Þ *=n* ¼ 0 , favorable 0ð Þ <1*=n* <1 , or unfavorable 1ð Þ *=n* >1 [43]. **Table 1** displays the correlation coefficient (*R*2 ) of these isotherms for Cr(III) on SDS-chitosan along with other relevant parameters.

*Adsorption of Chromium from an Aqueous Solution onto Chitosan Beads Modified with Sodium… DOI: http://dx.doi.org/10.5772/intechopen.104093*

**Figure 15.** *(a). Adsorption isotherms of Cr(III) using SDS-chitosan. (b). Adsorption isotherms of Cr(VI) using SDS-chitosan.*

#### **Table 1.**

*Isotherm parameters for Cr(III) and Cr(VI) adsorption onto SDS-chitosan.*

As shown in **Figure 16(a)**, **(b)**, and **Table 1**, it was revealed that the *R*2 values for Cr(III) and Cr(VI) were relatively large and a favorable adsorption of Cr(III) and Cr (VI) on SDS-chitosan was presented. Compared to the Freundlich isotherm, the Langmuir isotherm had a larger *R*2 value, which suggests that the adsorption of Cr(III)

**Figure 16.** *(a). Langmuir isotherm of Cr(III) adsorption onto SDS-chitosan. (b). Langmuir isotherm of Cr(VI) adsorption onto SDS-chitosan.*

and Cr(VI) on SDS-chitosan mainly occurred by a monolayer reaction. The maximum adsorption capacity (*qmax*) calculated from the Langmuir model was 3.42 mg/g for Cr (III) and 3.23 mg/g for Cr(VI).

#### **3.4 Kinetic studies**

The rate-controlling steps of the adsorption system are essential to survey the mechanism of Cr(III) and Cr(VI). Kinetic studies were conducted to explain the adsorption mechanism of Cr(III) and Cr(VI) ions onto the SDS-chitosan beads. The effects of contact time on the kinetics of Cr(III) and Cr(VI) adsorption by SDSchitosan adsorbent are displayed in **Figure 17(a)** and **(b)**. According to **Figure 17(a)**, the removal of Cr(III) from SDS-chitosan increased sharply in the initial 6 h, indicating that the uptake of Cr(III) was primarily caused by chemical sorption. As shown in *Adsorption of Chromium from an Aqueous Solution onto Chitosan Beads Modified with Sodium… DOI: http://dx.doi.org/10.5772/intechopen.104093*

**Figure 17.** *(a). Pseudo-first-order and pseudo-second-order Cr(III) adsorption on SDS-chitosan. (b). Pseudo-first-order and pseudo-second-order Cr(VI) adsorption on SDS-chitosan.*

**Figure 17(b)**, the adsorption of Cr(VI) by SDS-chitosan beads increased significantly within 96 hours, although the increase continued at a slower rate for 48 h. The rapid adsorption within the initial 24 h indicated that Cr(VI) uptake was mainly dominated by chemical sorption or surface complexation.

To investigate the mechanism of adsorption of Cr(III) and Cr(VI) on SDSchitosan, fitting was determined according to the pseudo-first-order (**Figure 18(a)** and **(b)**), pseudo-second-order (**Figure 19(a)** and **(b)**) and intraparticle diffusion kinetic models (**Figure 20(a)** and **(b)**). The equations of these models are given by:

Pseudo-first-order model:

$$
\ln \, = \ln \left( q\_e \right) - k\_1 t \tag{4}
$$

Pseudo-second-order model:

$$\frac{t}{q\_t} = \frac{1}{k\_2 q\_e^2} + \frac{t}{q\_e} \tag{5}$$

**Figure 18.** *(a). Pseudo-first-order linear kinetic model of Cr(III) adsorption by SDS-chitosan. (b). Pseudo-first-order linear kinetic model of Cr(VI) adsorption by SDS-chitosan.*

Intraparticle diffusion model:

$$q\_t = K\_{id} t^{1/2} \tag{6}$$

where qe and qt are the adsorption capacities of Cr using the SDS-chitosan beads at the equilibrium and time, *t*, respectively (mg∙g�1 ), *k*1is the rate constant of the pseudo-first-order adsorption (h�1 ), *k*2 is the rate constant of the pseudo-secondorder adsorption (g∙mg�1 ∙h�1 ) and *Kid* is the rate constant of intraparticle diffusion (mg∙g�1 ∙h�1/2) [44–46].