A review on nickel(II) adsorption in single and binary component systems and future path

Journal of Environmental Chemical Engineering 7 (2019) 103305

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A review on nickel(II) adsorption in single and binary component systems and future path

T

Md. Aminul Islama,b,⁎⁎, Md. Rabiul Awualc,d,⁎, Michael J. Angovea a

Department of Pharmacy and Biomedical Science, La Trobe Institute for Molecular Science (LIMS), La Trobe University, Bendigo, VIC-3550, Australia Department of Arts and Sciences, Faculty of Engineering, Ahsanullah University of Science and Technology (AUST), Dhaka, 1208, Bangladesh c Materials Science and Research Center, Japan Atomic Energy Agency (JAEA), Hyogo 679–5148, Japan d Center of Excellence for Advanced Materials Research, Faculty of Science, King Abdulaziz University, Jeddah 21589, Saudi Arabia b

ARTICLE INFO

ABSTRACT

Keywords: Adsorption Nickel(II) Wastewater Single and binary Surface complexation Future path

Water polluted with heavy-metal ion has been a major problem in recent years. Among various metal ions, nickel (II) is a priority pollutant commonly found in industrial wastewater. As a highly toxic element at an elevated concentration, Ni(II) can pose a serious threat to our ecological environment as well as human being. Ni(II) adsorption from wastewater is a must for environmental management and sustainability. Remediation of Ni(II) contaminated water is possible through adsorption onto various innovative adsorbents from the aquatic environment. The current review looks at the present status of the research done so far Ni(II) adsorption using various adsorbents from wastewater. Ni(II) adsorption kinetics, edges, isotherm, thermodynamic parameters, and Ni(II) adsorption mechanism have also been talked over. Efforts have also been made to steer out of the advantages and disadvantages of adsorbents and the future research need in Ni(II) adsorption by adsorbents. Agricultural based substrates and nanosized metal oxides have been found a hopeful alternative for Ni(II) adsorption from wastewater. The Ni(II) primarily adsorbed onto a homogeneous substrate forming a monolayer. Ni (II) generally formed outer-sphere complexes at low pH values while it formed inner-sphere complexes at higher pH. More than one species is being sorbed, or more than one type of surface site is involved in Ni(II) adsorption process or both. However, significant research is needed to understand Ni(II)-surface interaction mechanism at the solid-water interface. This review can fill the lacuna of researchers who would like to do more research in this related area in depth.

1. Introduction

are readily absorbed by the human skin [6]. Thus, it is mandatory to treat wastewater polluted with Ni(II) prior to their release into the environment. However, the regulation limit for Ni(II) in the water, air, soil, and food depend on many factors. A study conducted by Cempel et al. [7] mentioned that the regulation limit for Ni(II) in water, air, soil, and food were 20 μg/L, 0.025 μg m−3, 50 mg/kg d.w. and 132 μg /day respectively. Moreover, investigation of Ni(II) speciation at the liquid-surface interface is vital for an understanding of its ecological risks, toxicity and thus Ni(II) adsorption from the aquatic environment. Nickel (atomic number 28 and 24th most commonly found) is an element of environmental concern [8]. Ni(II) belongs to the suite of toxic metal ions (Co2+, Ni2+, Cd2+, Cu2+, Zn2+ Pb2+, Hg2+) which are subject to concentration limits in drinking water [8,9]. Wastewater polluted with metal ions must be treated carefully before its disposal, and its long-term behaviour in a geological repository or surface

Nickel (II) polluted water is a major problem in recent years due to its toxicity and tendency to bioaccumulate [1]. Due to its hazardous nature, Ni(II) has been recognized as priority pollutants that have found to deposit in the environment causing adverse effects to animal species [2]. The primary sources of Ni(II) are the industrial effluents from mining, oil refining, mineral processing, electroplating, forging, silver refining, paint formulation, battery manufacturer, and steam electric power plants [2,3]. While at trace levels Ni(II) is micronutrient, it is a toxic pollutant and influences animal and human health if excess Ni(II) is ingested [4]. Acute ingestion or intake of Ni(II) has severe health effects, including diarrhoea, renal oedema, gastrointestinal ache, pulmonary fibrosis, cardiovascular and kidney disease and cancer [1,5]. In addition, several Ni(II) compounds namely carbonyls are hazardous and

Corresponding author at: Materials Science and Research Center, Japan Atomic Energy Agency (JAEA), Hyogo 679–5148, Japan. Corresponding author at: Department of Pharmacy and Biomedical Science, La Trobe Institute for Molecular Science (LIMS), La Trobe University, Bendigo, VIC3550, Australia. E-mail addresses: [email protected] (Md. A. Islam), [email protected] (Md. R. Awual). ⁎

⁎⁎

https://doi.org/10.1016/j.jece.2019.103305 Received 21 May 2019; Received in revised form 25 June 2019; Accepted 19 July 2019 Available online 22 July 2019 2213-3437/ © 2019 Elsevier Ltd. All rights reserved.

Journal of Environmental Chemical Engineering 7 (2019) 103305

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disposal site is of high environmental relevance. In addition, large amounts of 59Ni (T1/2 = 76,000 years) and 63Ni (T1/2 = 100 years) occur as fission products in nuclear waste and thus the fate of these radionuclides has to be considered in performance assessment investigations of planned repositories for nuclear waste [8]. Based on the US Environmental Protection Agency (EPA) and the World Health Organisation (WHO), the highest satisfactory concentration for Ni(II) in drinking water is 0.02 mg/L [10]. Given extensive Ni(II) pollution and improper drinking water guideline, at that place is notable research interest to remove Ni(II) before release wastewater. To date, several procedures for the treatment of Ni(II) polluted water have been widely investigated. The typical traditional methods applied for metal ions from industrial wastewater, and soils include solvent extraction, adsorption, photocatalytic purification, microwave catalysis, reverse osmosis, ion-exchange, chemical precipitation, ultrafiltration, phytoextraction, and electrodialysis [11–29]. Each of the methods has its own advantages and disadvantages and detailed discussion on their comparative pros and cons have been reviewed elsewhere [28,30]. However, the selection of the procedure is based on cost involvement and concentration of wastewater. Among these methods, adsorption is now familiar as a promising and economically viable approach [31,32]. Adsorption process offers flexible design, high efficiency, and operation of treatment processes and producing sludge free effluent in many instances. Adsorption can be defined as a surface miracle in which adsorbents species are readily attracted and thus held to the surface functional group of a substrate at the sufficient equilibrium state [33–36]. Because of the compilation of recent research in environmental chemistry, some critical reviews relating to the adsorption of various pollutants have been published. Examples include Ni(II), Cd(II), Co(II), Mn(II), Hg(II), phosphate and heavy metal ions adsorption from wastewater using various adsorbents [37–44]. Raval et al. critically assessed the corresponding research on Ni(II) adsorption onto various innovative substrates such as activated carbon (AC) materials, low-cost substrates, nanomaterials, nanosized-metal oxides nano-composites, and matrix materials [10]. They analyzed each paper in terms of experimental parameters and maximum adsorption capacity. They found that agricultural based solid waste and biosorbents exhibited promising adsorption capabilities for Ni(II) ions from the water system. In our previous article, we have reviewed Co(II) adsorption from wastewater using various innovative adsorbents and their mechanism of adsorption [9]. In another review, Gikas et al. systematically discussed the influence of Ni(II) and Co(II) adsorption to the development of microorganisms grown under aerobic environment using activated sludge [45]. They found that these metal ions inhibited the growth of microbial at relatively high concentrations. Radenovic et al. summarised the related information on several features of low-cost materials as adsorbents, which derivative from agriculture, industry, and nature during Ni(II) removal [31]. To highlight their practical applicability, selected information on the key experimental aspects in terms of adsorption kinetic, edges, isotherm, and thermodynamic parameters have been presented. However, there is no systematic review for Ni(II) adsorption using diverse innovative adsorbents in single and binary solute systems together with their adsorption mechanism. The present review highlights the state-of-the-art research done so far relating to Ni(II) adsorption from an aquatic environment using both natural and synthetic materials designed for this purpose. The prime aim of this investigation is to summarize the Ni(II) adsorption mechanism in single and binary solute systems using various adsorbents. Several important experimental factors and reported maximum adsorption capacity values are provided. Special emphasis was also given to Ni(II) adsorption mechanism in terms of surface complexation model (SCM) and proton stoichiometry. Furthermore, the advantages of reusability of these adsorbents after remediation and future prospects of research in this area are discussed.

2. Nickel(II) adsorption onto adsorbents Interaction of metal ions with soil minerals is important for sustaining the environment and human health. Therefore, Ni(II) sorption has been investigated previously using a wide variety of surfaces [4,46–55]. It has been shown that Ni(II) sorption is highly dependent on contact time, solution pH, initial Ni(II) concentration, temperature, ionic strength, the existence of organic matter such as humic and fulvic acid, the nature and amount of adsorbent [47,49,51–55]. The Ni(II) adsorption depends on pH due to changes in both the metal speciation in solution and the nature of the surface [4,47,48,50,55–57]. Metal oxide minerals are more negatively charged with increasing pH due to the loss of protons by substrate functional groups and therefore, may sorb Ni(II) more readily from aqueous solutions [58]. The principal prerequisite for a viable Ni (II) adsorption process is the application of an adsorbent with high selectivity, high adsorption capability, and affordable and high regeneration capacity. AC has been a popular adsorbent for Ni(II) adsorption due to its high specific surface area (SSA), microporous character and the chemical nature of the substrates [31]. Nevertheless, due to the high cost of AC, there is intense interest in using other alternative substrates, including low-cost, metal oxides, clay minerals, zeolites, biopolymers, and miscellaneous adsorbents. The next segments will present Ni(II) adsorption using various adsorbents and their adsorption performance in terms of maximum surface coverage. 2.1. Activated carbon/carbonaceous substrates Carbonaceous substrates such as activated carbon (AC), carbon nanotubes, carbon fibers, and aerogels are promising substrates. Of them, AC has been widely practiced as a highly active adsorbent for adsorption of different pollutants, including Ni (II) from the aquatic environment. Carbonaceous materials are well developed SSA, high micropores volume, favorable pore size distribution and relatively higher adsorptive capacity and these are commonly used adsorbents for pollutants [31]. However, due to their high-cost involvement and regeneration cost, they are not viable for frequents use. Therefore, a large number of researchers are using and still looking for substitute adsorbents to adsorb Ni(II) ion from polluted water. Several researchers have studied Ni(II) removal using activated carbon and carbonaceous substrates [59–61]. Chen et al. demonstrated Ni(II) adsorption from simulated water using multi-walled carbon nanotubes (MWCTs, SSA 197 m2/g) and found that Ni(II) adsorption was highly dependent on solution pH and initial concentration of Ni(II) ion [62]. However, Ni(II) adsorption was not affected due to the change in the concentration of ionic strength. Pseudo-second order (PS2) and Langmuir equation were noted to fit experimental data well. Thermodynamic calculations confirmed Ni(II) adsorption to be spontaneous and endothermic. Furthermore, the ion-exchange mechanism was mainly accountable for Ni(II) adsorption onto the substrate. Salihi et al. prepared graphene oxide (GO) and sodium dodecyl sulphate (SDS) modified graphene oxide (SDS-GO) and were used these substrates to adsorb Ni(II) [61]. Ni(II) adsorption was noted to enhance noticeably from 20.2 to 55.2 mg/g observed by Langmuir model because of the functionalization of the substrate by SDS. Ni(II) adsorption was a purely electrostatic attraction, chemisorbed followed by ion-exchange. Abd ElMagieda et al. investigated the utilization of graphite activated carbon and modified activated carbon by Tetraethylenepentamine to adsorb Ni (II) from simulated water [60]. The observational data exhibited an endothermic and spontaneous process. Langmuir isotherm was able to describe the data well implying the formation of a monolayer onto a homogeneous substrate. A study of Ni(II) adsorption onto AC prepared from coir pith conducted by Kadirvelu et al. found that concentration of both Ni(II) and substrate played a vital role and an increase in Ni(II) concentration with fixed AC concentration leads to a better Ni(II) adsorption [63]. Both Langmuir and Freundlich equations were proposed 2

Journal of Environmental Chemical Engineering 7 (2019) 103305

Md. A. Islam, et al.

Fig. 1. The major mechanism involved in Ni(II) adsorption onto AC and AC-PER (Figure taken from Liu et al. [59]).

to fit Ni(II) adsorption data reasonably well. The maximum surface coverage was 62.5 mg/g at pH 5.0 and 30 °C. Ni(II) adsorption enhanced from pH 2.0 to 7.0 and then remained constant at pH 10.0. Liu et al. investigated Ni(II) adsorption using activated carbon prepared from lotus stalks and found that Ni(II) adsorption increased rapidly in the pH range of 2.5–4.0 thereafter increased slowly from pH 4.5 to 8.0 with increasing pH [59]. Ni(II) adsorption capacity of AC-PER was three times higher than that of AC indicating that surface chemistry played a vital role in the adsorption process. AC-PER contained more positively charged surface sites for the exchange with Ni(II) ion, a more stronger attraction with Ni(II) and more Ni(II)-surface complexation. On the basis of macroscopic and spectroscopic measurement (ATR-FTIR and XPS), they proposed that Ni(II) adsorption onto the substrates involved electrostatic attraction cation exchange, and surface complexation (Fig. 1). A list of the published papers [59–70] that used Ac and carbonaceous substrates as adsorbents involving experimental conditions together with maximum surface coverage to remove Ni(II) is shown in Table 1. It is seen that carbonaceous substrates are promising

adsorbents for Ni(II) removal. It is also remarkable that AC prepared from low-cost agricultural-based materials have been popular in recent years. Commonly, PS1, PS2, Langmuir, and Freundlich model has been useful for modeling Ni(II) adsorption data. In many instances, the Langmuir equation fitted experimental data well while PS2 described kinetic data satisfactorily. It is mentionable for future researchers that only a few studies involved Ni(II) adsorption using spectroscopic techniques and no attempts have been made to adopt surface complexation model(SCM) for Ni(II) adsorption data. This can be a promising future research area. 2.2. Agricultural based low-cost substrates Agricultural and natural waste-based substrates are the major sources of low-cost adsorbents [64–71]. The low-cost agricultural wastes have been promising bioadsorbents either in their original form or after some amendment since they are comparatively inexpensive and often obtainable in large masses as to the residues from agricultural actions. Ni(II) adsorption performance by several substrates have been

Table 1 Experimental conditions and maximum adsorption capacity (Qmax) of various carbon materials for Ni (II) removal. PS1 = pseudo-first-order kinetics; PS2 = Pseudosecond-order; IPD = Intraparticle diffusion; E = Elovich model; L = Langmuir; F = Freundlich; T = Temkin; D–R = Dubinin-Radushkevich model. Adsorbent

SBET (m2/ g)

Time (h)

pH

T(˚C)

[Ni(II)]0(mg/ L)

[Dose](g/ L)

(qmax) (mg/g)

Kinetic

Isotherm

References

Multi-walled carbon nanotubes (MWCTs) Activated carbon from coirpith

197 /

/ /

6.5 /2-7

18 30

2-20 30-50

0.75 0.2

8.77 62.5

L L

[62] [63]

Graphene oxide (GO) Sodium dodecyl suphate modified graphene oxide (GO) AC from Cucumis peel AC

/ /

24 24

3-9 3-9

25 25

40 40

0.1 0.1

20.19 55.16

PS1, PS2 PS1, PS2 PS1, PS2 PS1, PS2

L, F L, F

[61] [61]

/ /

/ 12

6 8

25 ± 2 40

100-400 10

2.5 0.5

/ 3.8

E L, F

[64] [65]

Polyvinyl alcohol/CNTs Nano porous composite (3DPCA) Activated carbon Activated carbon-pentaerythritol (AC-PER) Spent coffee-based AC Spent coffee husk-based AC Activated carbon prepared from almond husk AC based on a native lignocellulosic precursor Chemically AC residue from biomass (ACR) Carbon residue (CR) Commercial AC

/

2

/

/

400

/

225.6

PS1IPD PS1, PS2, IPD /

/

[66]

1419 343 383 464 / 1638 259 14.4 603

48 48 24 24 24 24 24 24 24

2.5-8.0 2.5-8.0 2-8 2-8 5 5.8 4-8 4-8 4-8

22 ± 1 25 ± 1 25 ± 1 25 ± 1 20 28 / / /

29.3 29.3 45 45 25 0.34 mM 75 75 75

0.4 0.4 0.2 0.2 0.12 0.1 0.125 0.125 0.125

34.04 54 57.14 51.91 30.77 0.13 mmol/g 18.2 5.6 3

/ / PS1, PS1, / PS2 PS1, PS1, PS1,

L, L, L, L, L, L L, L, L,

3

PS2 PS2 PS2, E PS2, E PS2, E

F F F, D-R F, D-R F F, D-R F, D-R F, D-R

[59] [59] [67] [67] [68] [[69] [70] [70] [70]

Journal of Environmental Chemical Engineering 7 (2019) 103305

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reported in previous works [72–74]. Compared to AC, agriculturalbased materials show better Ni(II) adsorption capacity. Saeed et al. found that black gram husk could adsorb heavy metal ions from simulated water follow the order of Pb(II) > Cd(II) >Zn(II) > Cu(II) > Ni(II) [74]. The maximum surface coverage of these metal ions was 50, 40, 34, 26 and 20 mg/g respectively using the substrate, and the adsorption phenomena were explained well by both the Langmuir and Freundlich equations. Chand et al. used apple pomace to adsorb Cd(II), Ni(II) and Pb(II) from simulated wastewater and the maximum surface coverages were 112, 51, and 179 mg/g, respectively [73]. Garba et al. explored Ni(II) and Cd(II) adsorption from aqueous solution onto modified plantain peel [72]. Under the optimized adsorption conditions, adsorption capabilities for Ni(II) and Cd(II) were 77.52 and 70.92 mg/g, respectively. Demirbas et al. explored Ni(II) adsorption from water using hazelnut shell activated carbon by a batch study [75]. The Langmuir equation correlated Ni(II) adsorption data onto AC nicely. They noted that Ni(II) adsorption enhanced with a growing temperature of the system as suggested by free energy change values (-5.58 to -7.01 kJ mol−1 at 20 °C to 40 °C) while Ni(II) adsorption onto the substrate was exothermic. Ramana et al. used pigeon peas hulls powder as an adsorbent to remove Pb(II) and Ni(II) from aqueous solution [76]. Both metal ions adsorption data followed the PS2 model while isotherm data described the Langmuir equation. Metal ion adsorption was spontaneous and feasible. Furthermore, the regeneration test was performed using HCl (0.05-0.25 M) as eluent and found that about 95% was recovered using 0.25 M HCl suggesting potential regeneration capacity of the surface. A list of papers [71,76,132,140,143,151–164] that used various low-cost materials as adsorbents to adsorb Ni(II) is displayed in Table 2. It is observed that many research groups have used low-cost adsorbents to remove Ni(II). This is due to locally available, cheap and environmentfriendly. In most cases, PS2 and Langmuir equation have been fitted to Ni(II) adsorption data reasonably well. As noted previously, no

attempts of spectroscopic and SCM have been adopted for Ni(II) adsorption onto low-cost substrates. 2.3. Metal (oxy-hydro) oxides In recent years, many researchers have used a large number of metal (oxy-hydro) oxides such as magnetite, goethite, and birnessite or their composites as adsorbents to remove Ni(II) from wastewater [77,78]. The ligand based materials are also used for the diverse toxic ion adsorption based on the sensitivity and selectivity [79–85]. However, the Fe, Mn, Al (oxy-hydro) oxides act as a natural filter of Ni(II) in water systems. These metal oxide materials act as hosts for metal ions with valences of +2, +3 and +5. The potential application of these materials and their composites for Ni(II) removal are reviewed below. Nanosized iron oxides have been used as adsorbents to adsorb Ni(II) from an aquatic environment as they have the feasibility of magnetic separation and regeneration aspects. Sharma et al. [54] explored Ni(II) adsorption onto Fe3O4 and found that Ni(II) sorption increased as the pH of the system increased. In addition, Ni(II) sorption was found to rise with increasing temperature of the system. Manganese oxides (MnOs) exhibit high surface areas (SSA) with notable microporosity and high affinity for metal ions [86–92]. Moreover, the redox nature of MnOs usually present as Mn(III, IV) affects the behaviour of redox-sensitive metal ion pollutants namely As(III/V), Se (IV/VI) and Cr(III/VI) [86,88,93]. Layer MnOs can adsorb a wide range of metal ions [9,86]. MnOs has a high potential for metal ion adsorption from wastewater. The use of MnOs as Ni(II) ion adsorbent has been explored in several previous studies [77,86,94–96]. Boonfueng et al. [96] examined Ni(II) and Pb(II) sequestration using hydrous MnO and MnO-coated montmorillonite. It was noted that both the metal ions adsorbed onto these substrates by forming inner-sphere complexes. Ni (II) coordinated to the vacancy sites in the MnO structure whereas Pb (II) adsorbed via bidentate corner sharing complexes. Ong et al. [95] used MnO2 to adsorb Ni(II) from simulated water and the Langmuir

Table 2 Ni(II) adsorption performance of various low-cost substrates under given experimental conditions. Adsorbent

SBET (m2/g)

Time (h)

pH

T(˚C)

[Ni(II)]0(mg/ L)

[Dose](g/ L)

(qmax) (mg/ g)

Kinetic

Isotherm

References

Green longan hull Rice husk ash Typha domingenesis biomass Green tea waste Rice husk ash (RHA) Carbon embedded silica from RHA

/ / / / 44 100

0.5 5 6 3 20 20

4.7 2-10 2-8 7 6 6

25 ± 1 30 25 ± 3 33 30-60 30-60

50 100 50 5 20-500 20-500

0.5 1 2.5 0.3 0.15-0.3 0.15-0.3

3.96 4.84 4.51 0.312 4.95 10.0

L, L, L, L, L, L,

[140] [132] [151] [143] [71] [71]

Raw red mud from aluminium industry Bagasse fly ash (BFA)

/

24

3-10

21 ± 1

58.7

0.1

27.43

PS2 / PS1 PS2 PS1 PS2 PS1, IPDPS2 PS1 PS2, IPD /

/

[152]

66.16

/

1-7

30

100-500

2

41.1

L, F

[153]

/ / / /

24 134 24 24

5.2 3 5-5.5 5-5.5

20 20 ± 2 / /

5-1000 29.34 29.34 29.34

0.1 0.1 2 2

15 1.11 5.2 8.0

PS1, PS2, IPD / / PS1, PS2 PS1, PS2

L, L, L, L,

F S F F

[154] [155] [156] [156]

415.82 / 0.62 / /

0.33 24 1 2 4 24

2-8 2-8 4 6 3-10 3-8

/ 20 ± 1 20 30 / 30

50 6 50-100 1-200 5-100 25-800

0.2 0.1 0.1 5 2.5 0.03-0.5

19.63 21.72 59 35.97 7.1 163.93

PS1, / PS1, PS1, / PS1,

L, L, L, L, L, L,

F, D-R F F F, D-R F F

[76] [157] [158] [159] [160] [161]

/ /

24 24

3-8 3-8

30 30

25-800 25-800

0.03-0.5 0.03-0.5

172.41 196.07

PS1, PS2 PS1, PS2

L, F L, F

[161] [161]

6.32

1

4

/

20

0.1

12.03

PS1, PS2

[162]

/

/ 2.5

8 2-5

25 30

50-500 50

10 0.5

44.47 9.10

/ PS1, PS2

L, F, T, D-R, Florry Huggins / L, F

Grape stalks waste Exhausted coffee waste Dry biomass of arthrospira pletensis Dry biomass of Chlorella vulgaris Pigeon peas hulls Red mud Lemna minor L Meranti sawdust Coal fly ash Cassia fistula leaves (golden shower) Cassia fistula stem bark Cassia fistula (golden shower) pods bark Jackfruit peel Rhizopus delemar Azadirachta indica leaves powder

4

PS2 PS2 PS2 PS2

F F, R-P F F F F

[163] [164]

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Md. A. Islam, et al.

model could describe Ni(II) adsorption data nicely. Ni(II) adsorption involved chemisorption by complexation of the functional group of MnO2. Ni(II) adsorption onto the substrate was spontaneous and thermodynamically favourable. Recently, Islam et al. [78] explored Ni(II) adsorption onto six different substrates such as birnessite, pyrolusite, hausmannite, manganite, boehmite and Mn-Al binary oxide from 1 mM NaNO3 electrolyte solution at 22 ± 2 °C. It was noted that Ni(II) had a much higher affinity for birnessite than for the other five minerals, while boehmite showed relatively little sorption (< 30%) below pH 8.5 where precipitation begins. The position of sorption edges (pH50) (derived from Ni adsorption edge curves) differed significantly and followed the order: birnessite (4.0) > Mn-Al binary oxide (5.6) > manganite (6.5) > pyrolusite (7.0) ˜ hausmannite (7.0) > Boehmite (9.0). Ni(II) adsorption isotherm data at pH 8.5 nicely described Langmuir equation implying that Ni(II) adsorbed onto these substrates by forming a monolayer. Birnessite showed a maximum adsorption capacity compared to other substrates. Apart from Fe and Mn oxides, several researchers also used aluminium oxides to adsorb Ni(II) from water [97,98]. Rajukar et al. [97] studied the possibility of activated alumina for Ni(II), Cr(III) and Cu(II) adsorption from synthetic solutions. The maximum adsorption capacities (mg/g) were in the order of Ni(II) (71.43) > Cu(II) (26.32) >Cr (III). Metal ions adsorption process followed Langmuir and Freundlich equations adequately while the process was thermodynamically spontaneous and endothermic. The positive value of entropy could be attributed to the rise in the randomness at the solid-water interface. Surface adsorption as well as ion-exchange well mainly responsible for these metal ions adsorption. Literature value of the maximum adsorption capacity [77,78,94,95,97,165–172] of some metal oxides is given in Table 3. Although previously pure metal oxides have been used, recently metal oxide composites have been popular for Ni(II) removal. In most cases, Langmuir and PS2 have found to describe Ni(II) adsorption data well.

2.4. Clay minerals Clay minerals are hydrous aluminosilicates formed from the colloid fractions of soil and water. Generally, they are comprised of the minerals such as metal oxides, carbonate and quartz [99]. The main clays include montmorillonite, bentonite, and kaolinite. Of them, montmorillonite shows maximum cation exchange performance and its price is twenty-times inexpensive than that of AC [99]. The negatively charged is neutralized by the uptake of positively charged species, resulting in clay the capability to attract and hold metal ions. The large SSA of clays also contributes to the high adsorption capacity. In addition, to improve the adsorption performance of the clay, it is treated with HCl. It is found that pre-treatment clay with HCl improved notably Ni(II), Cu(II) and Zn (II) adsorption. Acid treatment changed the surface chemistry of the clay, increasing its adsorption capacity. The maximum surface coverage follows the order of Cu(II)> Ni(II) >Zn(II) [100]. Several authors have used clay minerals in Ni(II)-polluted water treatment. Zhang et al. [101] used lignocellulose/montmorillonite nanocomposite to remove Ni(II) from the aquatic environment and Langmuir model was found to fit well yielding a maximum surface coverage of 94.86 mg/g at pH 6.8. PS2 was fitted to Ni(II) adsorption data. Using HNO3 as eluent the substrate was able to desorb 81.36 mg/g in 30 min. Yang et al. [58] explored Ni(II) adsorption onto bentonite surface and noticed that Ni(II) uptake was highly dependent on pH and ionic strength. Ni(II) sorption onto bentonite increased slowly with increasing pH from pH 2.0 to7.0, then increased abruptly with enhancing pH at around 8.0 and then leveled off at pH above 8.5. The rapid Ni(II) sorption at higher pH has been suggested to result from the formation of a precipitate of insoluble Ni(OH)2 on bentonite surface. Chantawong et al. [102] used kaolinite and ball clay as adsorbents to adsorb Cd(II), Ni(II), Cu(II), Zn(II), Pb(II) and Cr(III) from simulated solution. In case of kaolinite, metal ion adsorption followed the order of Cr(III)>Zn(II)>Cu(II) ≡ Cd(II)>Ni(III)>Pb(II) while that of for ball clay was Cr(III)> Zn(II) > Cu(II)>Cd(II)>Ni(III)>Pb(II). Ballclay exhibited higher metal adsorption than kaolinite because illite, the prime mineral of ball clay, has a higher surface positive charge than

Table 3 Ni(II) adsorption performance of various metal oxide substrates under given experimental conditions. Adsorbent

SBET (m2/ g)

Time (h)

pH

T(˚C)

[Ni(II)]0(mg/ L)

[Dose](g/ L)

(qmax) (mg/g)

Kinetic

Isotherm

References

Manganese dioxide δ-MnO2 Graphene nanosheet/d-MnO2 (GNS/ MnO2) composite Magnetic manganese dioxide Activated alumina Gibbsite

/ / /

24 2 2

4.5-7.5 7 7

25 25 25

10-200 50 50

10 / /

114.9 30.63 46.55

PS1, PS2 / /

L, F L L

[95] [94] [94]

240 / 13

3 1 24

4-7 7 4-9

27 25

0.08-1 5-50 100

5 0.5 5

/ / /

95

24

4-9

25

100

5

Cyclodextrin/MWCNTs/Iron oxides Fe3O4-MnO2

64 135.7

4 3

6.5 6

25 25

10 5-75

0.5 0.2

Polypeptide (γ-glutamic acid)-Fe3O4GO-(o-MWCNTs) Hydrous TiO2

/

1

2-10

25

190

0.4

384.62

663

1

2-10

15-45

10-150

0.1

22.7

L, F, R-P

[169]

Polyacrilamide/ Fe3O4 microcomposite

/

24

1-6

/

1-500

0.1

0.155 mmol/g

L, F, S

[170]

Fe3O4 -tea waste Zero valent iron nanoparticles Fe3O4- GS Birnessite Pyrolusite Hausmannite Manganite Boehmite Mn-Al binary oxide

27.5 / 62.43 363 4.8 55 66 220 /

2 2 0.5 0.5 0.5 0.5 0.5 0.5

6 1-7 2.5-12.5 2.5-12.5 2.5-12.5 2.5-12.5 2.5-12.5 2.5-12.5

30

50-100 20-150 5 ˜6 ˜6 ˜6 ˜6 ˜6 ˜6

0.25 1 0.1 0.28 20.8 1.8 1.5 0.45 0.73

38.3 133.3 158.5 4.67 μmol m−2 1.71 μmol m−2 4.12 μmol m−2 4.61 μmol m−2 3.22 μmol m−2 4.35 μmol m−2

PS1, PS2, IPD PS1, PS2, E, IPD PS1, PS2, IPD

L L, F L, F, R-P, Hill L, F, R-P, Hill L, F L, F, S, T, RP L, F

[77] [97] [165]

Goethite

35.3 71.43 9 × 10−5 mol/ m2 2 × 10−5 mol/ m2 38.24 55.63

L, L, L, L, L, L, L, L, L,

[171] [172] [78] [78] [78] [78] [78] [78]

20-40 22 ± 2 22 ± 2 22 ± 2 22 ± 2 22 ± 2 22 ± 2

5

/ / PS1, PS2

PS1, PS2, E PS1, PS2 / / / / / /

F F, D-R F F F F F F F

[165] [166] [167] [168]

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known structures, having exchangeable Na+, K+, Mg2+ and Ca2+ ions and water in their structural framework [105,106]. These ions can occupy the channels in the structure and are substituted with metal ions. The most common formula of a zeolite is sodium alumino orthosilicate (Na2O.Al2O3.xSiO2.yH2O). In nature, several classes of zeolites occur and are well distributed. Because of their unique surface properties, they are widely applied in the field of catalytic industry, material science and water purification [105,106]. Yang et al. [105] applied mordenite zeolite to adsorb Ni(II) from simulated water and noted that Ni(II) adsorption was depended on ionic strength below pH 7 while it was independent above pH 7.0. The EXAFS findings recommended that the 63Ni(II) adsorbed by ion-exchange leading to the development of outer-sphere complexes (Fig. 3, a). Initially, for rapid uptake, 63Ni(II) adsorption was dominated by the establishment of inner-sphere surface complexes (Fig. 3b). Upon prolongation, 63Ni(II) adsorbed by the development of Ni phyllosilicate co-precipitates and/or an insoluble precipitate of Ni(OH)2(s) (Fig. 3c). The second shell fit result indicated the probable establishment of Ni polynuclear surface complexes (binuclear dimers) (Fig. 3d). Irannajad et al. [106] removed the Co(II), Ni (II) and Pb(II) from water using manganese oxide coated zeolite. The isotherm data indicated that all metal ion sorption data followed both the Langmuir and Freundlich equation while kinetic data described PS2 model well. Thermodynamic revealed that Pb(II) was more spontaneous than those of Co(II) and Ni(II) ions. Application of several zeolite materials in Ni(II) adsorption [48,105,106,153] is shown in Table 5. Like previous tendencies, most Ni(II) adsorption data fitted well with a PS2 and Langmuir equation.

Fig. 2. Influence of pH on the percentage removal of Ni(II) and Mn(II) onto kaolinite clay. [Ni(II)]0 = [Mn(II)]0 = 100 mg/L, m = 0.1 mg/L, 27 °C (Figure adapted from Dawodu et al. [104]).

kaolinite surface. In another study, Yang et al. [103] used montmorillonite as an adsorbent to adsorb Cu(II) and Ni(II) from simulated water in monocompetitive systems. It was found that the presence of Ni(II) had no notable influence of Cu(II) adsorption while the presence of Cu(II) inhibited Ni(II) adsorption. Ni(II) shaped microstructure varied from the hydrated free Ni(II) ions at pH 5 and 6 while it formed inner-sphere complexes at pH 8.0 and above this pH Ni(II) formed Ni(OH)2 precipitation with the substrate. In another report, Dawodu et al. [104] used unmodified Nigerian Kaolinite clay to remove Ni(II) and Mn(II) using the batch method. They found that at low pH values of 2.0–4.0, Mn(II) ions were adsorbed more that Ni(II) but the opposite trend was observed at higher pH values (Fig. 2). Freundlich equation could fit isotherm data well implying heterogeneous surface nature whilst the PS2 model described kinetic data satisfactorily. The maximum surface coverages for Ni(II) and Mn(II) were 167 mg/g and 111 mg/g respectively at pH 6.0 and 27 °C. Thermodynamic data indicated that the removal of these metal ions was spontaneous, endothermic and physisorption process. Examples of clay minerals for Ni(II) adsorption [46,101,104,173–177] are represented in Table 4. It is noted that claycomposites have been commonly using as an adsorbent to adsorb Ni(II) from simulated wastewater.

2.6. Biopolymer-based adsorbents Biopolymers are typically comprised of lignin, chitin/chitosan, and cellulose as major ingredients with few hydroxyl and carboxyl groups which have the capability to bind metal ions by donation of an electron pair to form surface complexes [107]. These adsorbents can improve mechanical strength and better resistance to chemical conditions after modification [107]. Recently, biopolymer-based substrates have attracted the attention of researchers for their application in water treatment because of their abundance and renewable nature. Irani et al. [108] prepared functionalized polyvinyl alcohol/tetraethyl orthosilicate hybrid membrane by sol-gel method and used this membrane to remove Cu(II) and Ni(II) from aqueous solutions. It was found that metal ions adsorption enhanced with growing pH and reached maximum at pH 5. Adsorption was described by PS2 and Freundlich model and was thermodynamically endothermic in nature. Mende et al. [109] used chitosan to adsorb Ni(II) from simulated water and found the maximum removal capacities were 82 and 21 mg/g respectively for NiSO4 and Ni(NO3)2 at pH 6.1. The thermodynamic study indicates that Ni(II) adsorption onto the substrate was spontaneous and

2.5. Zeolite minerals In the distant past, zeolites attained notable attention because of their ion-exchange capacity to adsorb metal ions from water. Zeolites are generally microporous, crystalline aluminosilicate substrates with

Table 4 Ni(II) adsorption performance of various clay substrates under given experimental conditions. Adsorbent

SBET (m2/ g)

Time (h)

pH

T(˚C)

[Ni(II)]0(mg/ L)

[Dose](g/ L)

(qmax) (mg/ g)

Kinetic

Isotherm

References

Lignocellulose/montmorillonite nanocomposite Nigerian kaolinite clay Pyrophyllite Montmorillonite Epichlorohydrin crosslinked chitosan-clay beads E. Coli biofilm supported on kaolin Calcium bentonite Un-calcined sodium exchanged (NaMMT) Modified montmorillonite (A-MMT)

/

1

6.8

60

18.8

0.1

94.86

PS1, PS2

L, F

[101]

13.41 / 32 /

3 24 24 24

2-8 5-8.5 3-11 6-7

27 25 / 20-45

100 58.7 50-500 100

0.1 20 0.5 0.1

167 / 100% 32.36

PS1, PS2, E / / /

L, F, T, D-R / / L, F

[104] [46] [173] [174]

/ 41.7 285.6

240 168 /

/ 7 2.3-9.8

37 25 25

10-100 / 50-100

1 0.2 6

/ 0.581 9.93

/ / PS1, PS2, E

[175] [176] [177]

190

/

2.3-9.8

25

50-100

6

7.78

PS1, PS2, E

F, S, R-P L, F, R-P L, F, R-P, Florry Huggins L, F, R-P, Florry Huggins

6

[177]

Journal of Environmental Chemical Engineering 7 (2019) 103305

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Fig. 3. Probable Ni(II) adsorption mechanism onto mordenite substrates: (a) Outer-sphere surface complexation/ ion exchange; (b) inner-sphere surface complexation; (c) (Co-) precipitation; (d) surface dimerization (Figure adapted from Yang et al. [105]).

endothermic. In addition, the role of the valence of the oxyanion on adsorption capacity of substrate for Ni(II) adsorption has proposed in Fig. 4. They found that higher anion valence (SO42−) led to a better Ni (II) adsorption capacity because of more electrostatic interactions between opposite ions in solution and adsorbed species even after monolayer formation. Table 6 summarises the studies of Ni(II) adsorption using biopolymer adsorbents [108,114,116,178–191]. It is seen that a wide range of biopolymers has been used as an adsorbent to remove Ni(II) from the water. To improve adsorption capacity, many researchers modified biopolymers which could transfer mass and exposed active bonding sites by coating or cross-linking.

single solute system, Ni(II) adsorption described Freundlich equation while Cr(VI) adsorption fitted Langmuir equation well. Maximum Ni(II) adsorption was 2.9 mg/g at pH 3.0 whilst it was 3.64 mg/g for Cr(VI) at pH 8.0. Adsorption data for both metal ions followed the PS2 model. The presence of Ni(II) reduced the amount of Cr(VI) adsorption while the presence of Cr(VI) enhanced the Ni(II) adsorption. In addition, an increase of temperature caused the increase of metal ions adsorption and this has been described to the acceleration of the diffusion of metals from the external layer into micropores of the substrate and adsorption process was spontaneous, endothermic and rise in the degree of freedom at the solid-water interface. Torab-Mostaedi et al. [110] used expanded perlite as adsorbents to sorb Cd(II) and Ni(II) from aqueous solution. The maximum removal efficiency of Cd(II) and Ni(II) was 88.8% and 93.3% at pH 6.0 using the amount of perlite 10 g/L and 8 g/ L respectively. Freundlich equation was found to fit equilibrium isotherm data adequately than the Langmuir one. PS2 model described kinetic data well. Both the metal ions adsorption was exothermic and feasible in nature.

2.7. Miscellaneous adsorbents Apart from the above-mentioned adsorbents, several authors also used various miscellaneous adsorbents to adsorb Ni(II) from aqueous solutions. Sha et al. [25] explored Cr(VI) and Ni(II) adsorption using diatomite waste modified by EDTA (SSA 9.86 m2/g; pHPZC 5.39). In a

Table 5 Ni(II) adsorption performance of various zeolite substrates under given experimental conditions. Adsorbent

SBET (m2/g)

Time (h)

pH

T(°C)

[Ni(II)]0(mg/L)

[Dose](g/L)

(qmax) (mg/g)

Kinetic

Isotherm

References

Supported zeolite-Y hollow fiber membranes

/

2

5

RT

10

0.02

126.2

L, F

[48]

Mordenite Manganese oxide coated zeolite Conventional zeolitic bagasse fly ash Electrolyte treated zeolitic bagasse fly ash

/ / 225.52 274.12

24 0.5 /

2.5-10.5 3-8 1-7 1-7

20 20 30 30

5-100 100 100-500 100-500

0.5 1 2 2

100% 7.9 73.85 93.72

PS1, PS2 / PS1, PS2, IPD PS1, PS2, IPD PS1, PS2, IPD

/ L, F L, F L, F

[105] [106] [153] [153]

7

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Fig. 4. Ni(II) adsorption mechanism onto chitosan substrate and Ni(II) sourced from (a) NiSO4 solution and (b) Ni(NO3)2 solution (Figure taken from Mende et al. [109]). Table 6 Ni(II) adsorption performance of various biopolymer substrates under given experimental conditions. Adsorbent

SBET (m2/ g)

Time (h)

pH

T(˚C)

[Ni(II)]0(mg/L)

[Dose](g/ L)

(qmax) (mg/ g)

Kinetic

Isotherm

References

Chitosan Graft copolymer of cellulose Starch-Graft-Acrylic Acid Copolymers Modified carboxymethyl cellulose hydrogel Carboxylated sugar bagasse Unmodified modified hydrochar Acid-hydrochar Alkali-hydrochar Chitosan Chitosan/nanoclay composite EDTA-Chitosan

/ / / / / / / / / 14.55 0.71

0.5 6 8 10 1 24 24 24 1 2 24

5.2 5 2-6 5 5.75 5.5 5.5 5.5 6 / 2.1

30 30 / 30 25 / / / 25 25 /

100 200 293.45 10 46.36 50 50 50 100 58.7 20-100

1 1-5 0.5 0.1 0.2 0.5 0.5 0.5 5 0.3 2

315 1770 14.7 1000 1.53 20.27 25.5 29.0 52.6 65 71.0

PS2,

L, L, / L, L, L, L, L, L, L L,

F, S

[178] [179] [180] [181] [182] [183] [183] [183] [184] [185] [114]

DTPA-Chitosan

0.36

24

2.1

/

20-100

2

53.1

PS2,

L, F, S

[114]

PVA/TEOS/TMPTMS hybrid membrane Histidine modified chitosan beads Grafted hydrazinylamine magnetite-chitosan Chitosan Ca-alginate Chitosan chelating resin Natural wool chelating fibers Thiourea modified chitosan Cross linked metal imprinted chitosanepichlorohydrin

182.3 / / / / / / 62.3 /

4 24 24 24 24 8 3 8 4

2-7 4-9 5 2-10 2-10 1-7 1-7 1-6 5

25-45 25 25 25 25 28 15-35 28 25

100 25-1000 120 6 6 100 100 100 10

0.1 0.1 0.25 0.5 0.5 0.3 0.2 0.3 0.1

10.29 140.8 254.12 41 179 40.15 46.7 15.3 29.23

PS1, PS1 / / PS1, PS1, PS1, PS1, PS1, / PS1, IPD PS1, IPD PS1, / IPD IPD IPD PS1, PS1, PS1, PS1,

PS2

L, L, L, L L L, L, L, L,

[108] [116] [186] [187] [187] [188] [189] [190] [191]

Recently, using trisodium citrate as crystal modifier are promising for metal ion adsorption. Jiang et al. [111] prepared flower like globular Mg(OH)2 and used this novel substrate to adsorb Ni(II) from simulated water. PS2 and Langmuir model were more suitable to describe Ni(II) adsorption data. Maximum surface coverage was 287.11 mg/g at 25 °C. Spectroscopic analyses (SEM, ATR-FTIR, and XPS) indicated that Ni(II) adsorption involved chemisorption rather than physisorption onto the substrate. The chemisorption process of Ni(II) removal has been proposed to occur as follows:

PS2

PS2 PS2 PS2 PS2 PS2

PS2 PS2 PS2 PS2

F F F F, S F F F F

F, D-R F F, S F, F, F, F,

T T T D-R

3. Comparative Ni(II) adsorption study Several authors explored the comparative investigation of Ni(II) adsorption using various adsorbents. Green-Pedersen et al. [56] undertook a comparative investigation of Ni(II) adsorption using MnO2, Fe(OH)3, montmorillonite, humic acid (HA), and calcite. Ni(II) adsorption data onto MnO2 was fitted by the Langmuir equation, whereas data for Fe(OH)3 were fitted satisfactorily by the Redlich-Peterson (R-P) model. Increasing ionic strength decreased Ni(II) sorption onto MnO2 and montmorillonite, whereas sorption was largely unaffected by ionic strength for Fe(OH)3. Results showed that the Ni(II) adsorption selectivity was in the order of MnO2 > Fe(OH)3 > humic acid ˜ montmorillonite > calcite. In another report Jacob et al. [112] examined Ni(II) adsorption from simulated water using activated carbon (AC), magnetosome, calcite, and magnetic calcite and noted that the maximum Ni(II) adsorption at pH 8.0 follows the order: calcite = magnetic calcite (100 mg/g) > activated carbon (3.7 mg/g) > magnetosome (2.1 mg/g). In all cases, Ni (II) adsorption was spontaneous, endothermic and improved randomness at the solid-water interface. Totlani et al. [71] demonstrated a

Ni2+ + Mg(OH)2 → Ni(OH)2 + Mg2+ Ni(OH)2 - e− + OH− → NiO(OH) + H2O Application of several miscellaneous adsorbents [29,110–117,192–198] for Ni(II) adsorption together with experimental situations is shown in Table 7. It is observed that a diverse range of adsorbents has been used as adsorbents to remove Ni(II) from aqueous solutions. A PS2 and Langmuir model could describe Ni(II) adsorption data nicely in most cases.

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Table 7 Ni(II) adsorption performance of various miscellaneous substrates under given experimental conditions. Adsorbent

SBET (m2/g)

Time (h)

pH

T(˚C)

[Ni(II)]0(mg/ L)

[Dose](g/ L)

(qmax) (mg/ g)

Kinetic

Isotherm

References

Expanded perlite Diatomite waste modified by EDTA calcium carbonate coated bacterial magnetosomes Calcite Magnetic calcite Dried activated sludge

1.89 8.86 /

2 3 12

6 2-10 8

20-50 30-50 40

10 40 10

15 0.1-0.6 0.5

2.24 6.52 2.1

PS1, PS2 PS1, PS2, IPD PS1, PS2, IPD

L, F L, F, D-R L, F, D-R

[110] [29] [112]

/ / /

12 12 4

8 8 2-7

40 40

10 10 5

0.5 0.5 0.5

100 100 2.37

PS2, IPD PS2, IPD PS2, E,

L, F, D-R L, F, D-R L, F, BET

[112] [112] [192]

Mesoporous nanosheets MgO Silica aerogel Amine modified silica aerogel Polyaniline/CoFeC6N6 nanocomposite Calcium hydroxyapatite Alginate -polyethylenimine (A-PEI) Alginate/Fucus- polyethylenimine (AF-PEI) Amino functionalised mesoporous silica Vermiculite-based zirconium oxides L-methionine modified Dowex-50 ion exchanger Vanadium mine tailing Flower globular magnesium hydroxide

101 1401.2 / 90.3 95 / / / 130.86 2.37

4 1 1 4 2 72 72 2 6 2

2-6 4 4 / 6.6 1-5 1-5 1.5-5.0 2-8 2-6

/ / / 25 20-70 20 20 25 25

500

2218 17.48 40.32 18.34 46.17 48.71 61 12.36 90.21 83.33

PS2

10 10-100 40-80 17.6 17.6 10-70 1-50 50-200

0.1 0.1 0.1 3.5 0.8 1 1 0.5-2 0.1 0.05

PS1, PS1, PS1, IPD PS1, / / PS1, / PS1, PS1, PS1, PS2 PS1,

L, L L L, L, L, L, L, L, L,

[193] [194] [194] [115] [195] [113] [113] [196] [197] [117]

42.39 101.5

24 2

2.5-8.0 2.14-7.71

30 10-45

/ 80

4 0.45

3.64 287.11

PS1, PS2 PS1, PS2, IPD

comparative study of Ni(II) removal using rice husk ash (RHA) and carbon embedded silica (CES) from RHA by batch technique. It was found that CES exhibited 18 folds higher Ni(II) adsorption performance over RHA. In both cases, Ni(II) adsorption followed the PS2 model. Langmuir equation displayed better fitting. They speculated that Ni(II) adsorbed onto this substrate by a reaction involving the exchange of cation via substitution of a proton from silanol group (≡SiOH) onto the surface. Further, they suggested that Ni(II) adsorption onto RHA could also take place by complex development after interaction between Ni (II) and substrate functional groups of both surfaces.

PS2 PS2 PS2 PS2 PS2, IPD

F F F, D-R F, S F, S F F F

L, F L, F

[198] [111]

5. Competitive adsorption measuring Ni(II) The ion selectivity is the key parameters for the specific metal ion adsorption using functional adsorbent materials [119–125]. Most laboratory-based investigation focusing on Ni(II) adsorption in the binary solute system. Moreover, multiple pollutants can occur in the environment which can change Ni(II) adsorption process. Previously several papers have investigated the competitive adsorption of Ni(II) to substrates in existence of a metal ion, an organic ligand, or a humic substance [119–131]. The existence of other metal ions or other sorbates has been found to enhance, inhibit, or have little effect on Ni(II) sorption depending on experimental conditions [50,127,132]. Strathmann et al. [50] studied Ni(II) adsorption onto boehmite (γ-AlOOH) in presence of fulvic acid (FA) and noticed that FA with an initial concentration of 10 mg/L did not markedly mark the sorption of 10 μM Ni (II), but 50 mg/L FA dramatically increased Ni(II) sorption below pH 7.5, with a slight reduction above pH 7.5. In a single solute system, Ni (II) developed a mononuclear bidentate edge-sharing surface complex with the substrate. Yang et al. [58] studied Ni(II) adsorption in existence HA onto bentonite and observed that presence of HA enhanced Ni(II) sorption below pH 8.0, while Ni(II) uptake was decreased at pH values beyond 8.0. The rise in Ni(II) sorption in presence HA at low pH has been ascribed to a reduction in the positive surface charge due to the accumulation of negatively charged HA on the substrate, which allowed electrostatic attraction between the surface and Ni(II) forming ternary Ni-HA-bentonite complexes. The decrease of Ni(II) sorption at higher pH has been suggested to be due to the development of a strong soluble Ni-HA complex in solution. In another study by the same author Yang et al. [133] investigated arsenic(III) and Ni(II) sorption in both single and binary component systems onto modified green tea waste. They observed that in the competitive sorption systems, the presence of both As(III) and Ni(II) caused the sorption of each to be inhibited by the presence of the other. They proposed that, as the amount of solid was constant, As(III) and Ni(II) competed for the limited number of surface sites and as a result exerted an antagonism effect on the sorption of one another. In a recent report, Sha et al. [134] noted that the existence of Cr(VI) enhanced the amount of Ni(II) adsorbed onto diatomite modified by EDTA, while the existence of Ni(II) reduced Cr(VI) sorption and this has been speculated to the availability of limiting binding surface sites.

4. Desorption and reusability studies Typically, an ideal substrate for Ni(II) adsorption should have a high SSA, porosity, better adsorption capacity, and be simply regenerated [113,114]. The successive regeneration and reuse of adsorbents are notable attributes from an economic and environmental viewpoint. Some investigations have also shown regeneration of the used adsorbents after Ni(II) adsorption [113–116]. Desorption experiments are primarily conducted by means of diluted acids such as NaOH, HCl or HNO3 [115]. Moazezi et al. [115] investigated regeneration test after Ni (II) adsorption using HCl, NaCl, and NaOH and found that using NaOH the regeneration capacity of the substrate were 98% (unmodified hydrochar), 99% (acid hydrochar), 91% (alkali-hydrochar) respectively. Demey et al. [113] investigated recycling capability of spent Alginate- polyethyleneimine (AF-PEI) and Alginate/Fucus- polyethyleneimine (AF-PEI) after adsorption of Cd(II), Pb(II), Cu(II), Ni(II) and Zn(II) using 0.1 M HCl/0.05 M CaCl2 solution. The substrates showed higher selectivity for Pb(II) and Cu(II) with a removal efficiency of more than 70% and 40% in five consecutive adsorption-desorption cycles. Kumar et al. [117] conducted regeneration study of L-methionine modified Dowex-50 ion exchanger using water (DW), 0.10 M HCl and NaOH solution. It was noted that 93% Ni(II) was recovered by the substrate in 0.1 M HCl solution. After the third cycle, 88% Ni(II) was recovered implying that the exchanger exhibited the higher regeneration capability. In another paper, Fu et al. [118] investigated the regeneration of spent substrates using 0.1 M NaOH followed by 0.01 M HCl and desorption rate after five consecutive cycles for Ni(II) adsorption was about 90% indicating high regeneration capacity.

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6. Ni(II) adsorption mechanism

described by empirical and SCM. Commonly applied isotherm models for describing the adsorption of Ni(II) include Langmuir, Freundlich, Sips, Temkin, Dubinin-Radushkevich (D–R), Redlich-Peterson and Toth model and detailed explanation of these model are provided in our previous report [9]. Langmuir equation assumes that only a saturation monolayer is formed on the homogenous surface of the adsorbent during adsorption. On the other hand, Freundlich equation is an empirical term based on Ni(II) adsorption on a heterogeneous surface [27,114]. Both the equations can be applied for the equilibrium adsorption data of many substrates implying that either monolayer or multilayer adsorption could take place on the substrate, based on the nature of the substrates. Application of various isotherm models is shown in Table 1–7 in the corresponding section. It was noted from the review that most of the Ni(II) adsorption data are best described by Langmuir isotherm.

In recent years, Ni(II) adsorption onto mineral surfaces has been widely investigated using macroscopic and spectroscopic techniques [46,47,49,50,105,135]. It has been proposed that Ni(II) adsorption mechanisms involve primarily cation exchange, surface complexation such as inner and outer-sphere complexation, surface tempted precipitation, coprecipitation, and diffusion [47,49,50,135,136]. Although macroscopic tactics are important to understanding Ni(II) adsorption in the environment, spectroscopic techniques including, XPS, XAS, EXAFS, and ATR-FT-IR are important methods to elucidate the Ni(II) adsorption process at the molecular level [105,135,136]. A combination of macroscopic data with spectroscopic analysis showed that Ni(II) sorption was mainly outer-sphere complexation and ion exchange at relatively low pH, while inner-sphere complexation and surface (co-) precipitation were of increasing importance at higher pH values [1,47,135]. Strathmann et al. [50] investigated the impact of fulvic acid (FA) on Ni(II) adsorption by boehmite using batch adsorption experiments and EXAFS spectroscopy measurements. They found that Ni(II) adsorption was in the development of inner-sphere bidentate complexes with aluminol groups (≡Al−OH) on the boehmite (γ-AlOOH) rather than the anticipated Ni-Al layered double hydroxide phase. The existence of FA increased the amount of Ni(II) sorbed by both ligand-bridged ternary Ni(II)-FA-boehmite complexes and metal-bridged ternary FA-Ni (II)-boehmite complexes. Recently, Qiang et al. [49] studied Ni(II) adsorption onto a calcareous aridisol (CA) soil using batch and spectroscopy techniques. They suggested that the sorption mechanism of Ni(II) on CA soil involved ion-exchange and inner-sphere surface complexes at low pH, while at higher pH the major mechanism involved the formation of Ni(OH)2 and a Ni-Al layered double hydroxide. Xu et al. [135] developed a mechanistic model for Ni(II) adsorption onto hydrous ferric hydroxide (HFO) using adsorption edge and EXAFS analysis and found that Ni(II) formed inner-sphere mononuclear bidentate edgesharing complexes with distances of Ni-O and Ni-Fe were 2.05–2.07 Å and 3.07–3.11 Å respectively. In another study, Sheng et al. [137] studied the Ni(II) adsorption mechanism onto diatomite using macroscopic batch and spectroscopic (XPS and EXAFS) analysis with the results showing that outer-sphere complexation dominated below pH 7, whereas inner-sphere complexes dominated in the pH range from 7.0 to 8.0, and surface coprecipitates dominated Ni(II) uptake at pH 10.0. They also found that a rise in the temperature resulted in a rise in the surface loading with surface coprecipitates dominating at higher temperatures.

6.3. Surface complexation modeling of Ni(II) adsorption Recently, surface complexation models (SCMs) have been simulated to understand the Ni(II) sorption mechanism. Several different surface complexation models have been proposed for Ni(II) sorption onto mineral surfaces [53,56,78,131,138–140]. Examples include use of the constant capacitance model (CCM) [47,131], extended constant capacitance model (ECCM) [78,139], diffuse layer model (DLM) [56,141] and triple layer model (TLM) [138]. Marmier et al. [139] modeled the sorption of Yb(III), Ni(II) and Cs(I) onto magnetite using the CCM, DLM and a non-electrostatic model (NOM) and remarked that the results of fits indicated the same stoichiometry for the surface reactions for each of these models order similar surface loading situations with the surface species [≡SOH2+], [≡SO−] and Ni(II) sorption species [≡SONi (OH)2−] successfully describing sorption for all of these models. Ding et al. [131] used both spectroscopic (FT-IR, EXAFS, and XPS) measurements and constant capacitance modeling (CCM) for simultaneous Pb(II), Cu(II) and Ni(II) adsorption onto carbonaceous nanofibers (CNFs) and found that inner-sphere surface complexation dominated adsorption of these metal ions. Gu et al. [138] reported the adsorption of Cd(II), Cu(II), Ni(II), Pb(II) and Zn(II) onto montmorillonite using the CCM, ECCM, DLM, and TLM at 25 °C. They proposed that metal ions were bound to the permanently charged basal surface sites of the substrate through the formation of outer-sphere complexes at low pH, while sorption occurred mainly as inner-sphere complexes on the variable charged edge sites at higher pH region. A similar trend has been also noticed by other groups for these metal ions adsorption onto kaolinite [140], and for Ni(II) sorption onto illite [47] and carbonaceous nanofibers [131]. Recently, Islam et al. [78] simulated Ni(II) adsorption edge and isotherm data onto six different substrates such as such as birnessite, pyrolusite, hausmannite, manganite, boehmite and Mn-Al binary oxide from NaNO3 solution by the extended constant capacitance model (ECCM) considering two inner-sphere complexes, [≡SONi+], [≡SONiOH] and an outer sphere complex, [≡SOH… NiNO3+]. The distribution of the sorbed species indicated that the main adsorbed species was [≡MnOH…NiNO3+] below pH 8.0, while at higher pH values the dominant species are [≡MnONi+], and [≡MnONiOH]. It was also noted that Ni(II) sorption isotherms at pH 6.0 onto mineral surfaces in the single solute system could be satisfactorily fitted using these ECCMs but the isotherm at pH 8.5 was not well fitted, possibly due to precipitation of Ni(OH)2 especially at higher concentrations.

6.1. Ni(II) adsorption kinetic modelling Ni(II) adsorption kinetic model is used to explore the rate controlling steps. Kinetic studies are executed in a way to vary the initial concentration of adsorbate, temperature, amount of adsorbent dose, pH, type and nature of adsorbents [114]. It is a continuous measurement of experimental data until sufficient equilibrium is achieved. The result is verified using kinetic equations to estimate the best fit, which will give a better understanding of the adsorption mechanism. The commonly used kinetic models for Ni(II) adsorption data are pseudofirst-order (PS1), Pseudo-second-order (PS2), Intraparticle diffusion (IPD) and Elovich (E) [114]. The detailed explanation of each of the model is provided in our previous article [9]. Utilization of different kinetics equations has been displayed in Table 1–7 in the corresponding section. From the literature search, it was evident that most of Ni(II) adsorption data followed the PS2 model well.

6.4. Proton stoichiometry for Ni(II) adsorption

6.2. Ni(II) adsorption isotherm modelling

Proton stoichiometry (χ), or proton exchange ratio refers to the moles of protons released for every mol of Ni(II) adsorbed onto a particular mineral surface [78]. The proton stoichiometry can be a useful parameter for predicting the type of reaction that might occur. Several authors have estimated the proton stoichiometry (χ) for sorption of

Generally, an adsorption isotherm can show a relationship between the amount Ni(II) adsorbed and Ni(II) concentration at a constant temperature at equilibrium [27,114]. Adsorption reactions are typically 10

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metal ions to mineral surfaces with values typically between one and two per cation sorbed, depending on the sorbing metal ion and the solution conditions [57,142,143]. Some authors have interpreted proton stoichiometries in terms of ion exchange reactions, where an ion with a 2+ charge might displace 2 protons to maintain charge balance [144]. More recently, higher stoichiometries have been ascribed to surface complexation reactions involving either multiple surface sites or surface hydrolysis [145,146]. Benjamin et al. [147] investigated the multi-site sorption of Cu(II), Cd(II), Zn(II) and Pb(II) onto amorphous iron oxyhydroxide and observed that the values of proton stoichiometry varied from 1.3 to 3.3 depending on the metal ion but was not significantly affected pH. Rodda et al. [148] investigated Pb(II) and Zn(II) sorption onto goethite and found the variation of χ between 1.0 and 2.0 with pH, temperature, and metal ions studied. Weerasooriya et al. [149] found a proton exchange ratio between 1.32 to 1.55 for the Cd (II)-gibbsite system. Proton stoichiometry for Co(II) sorption onto goethite was 2.3, while that of for Cu(II), Pb(II), Cd(II) and Zn(II) were 1.8, 2.4, 2.2 and 2.2, respectively [144]. Different values of (χ) for Ni (II) adsorption onto various mineral surfaces that suggests sorption of Ni(II) reactions for different substrates are notably different. This difference may be due to the difference in the crystal structure, specific surface area (SSA), and point of zero charge (pHPZC) in each mineral [148,150]. Thus, the finding of this study suggests that more than one species is being sorbed, or more than one type of surface site is involved in Ni(II) adsorption process or both [78].

i Execute more experiments on competitive Ni(II) adsorption in the presence of another metal ion, anion, dye or humic substance onto various adsorbents to understand to Ni(II) adsorption in multiple pollutant systems. ii Most papers described herein involved Ni(II) adsorption by batch mode and column study and only a few researchers dealt with a pilot or industrial scale. iii Reuse and proper management of spent adsorbents is an important topic and has not considered entirely. iv There is a bigger opportunity to develop toxic free and aggregation free nanoadsorbents to mitigate Ni(II) ion from wastewater. v More effective, efficient, reusable and eco-friendly materials should be developed, and experiments are carried out with real industrial effluents. vi Researches are necessary to have a comparative investigation of diverse adsorbents for Ni(II) adsorption from water taking account of the feasibility, cost-effectiveness under similar conditions. vii In recent years, although surface complexation models (SCMs) have proved promising tools to identify probable Ni(II) adsorption mechanism and only a limited number of papers have attempted this SCMs approach and poorly understood and thus more extensive research is needed to understand Ni(II) adsorption mechanism. viii More spectroscopic studies using zeta potential measurement, attenuated total reflection Fourier transform infra-red, X-ray photoelectron spectroscopy, extended X-ray absorption fine structure and X-ray absorption near edge structure are necessary to understand actual Ni(II) adsorption mechanism occurs at the solid-water interface. ix Recently, the application of composite nanomaterials in Ni(II) removal has increased considerably. Adsorption capacity is found to depend on various parameters. Moreover, there are various aspects associated with adsorbents which require detailed investigations. Attempts are also desirable to optimize the process parameters and laboratory scale experiments entail commercialization. x Little work is done on the use of recycling materials such as pulp and paper waste, garments, feedstocks, sludges, rubbers, plastics and medical wastes as adsorbents to remove Ni(II). This can be potential future research. xi Future extensive research is needed to use green waste such as leaves, barks, flowers, straws, and grass as adsorbents to adsorb Ni (II) on a large scale. xii Future research includes the utilization of animal skeleton, bone, skin, hair, feathers as adsorbents in Ni(II) adsorption study. xiii Compare Ni(II) adsorption performance with other removal techniques to understand the suitability of the process. xiv Recently, the application of density functional theory (DFT) has been very popular in metal ion adsorption study to estimate adsorption geometry, energy, and nature of the binding energy. However, this study was not conducted frequently for Ni(II) adsorption and this is an interesting research area.

7. Thermodynamic calculation for Ni(II) adsorption Temperature is a crucial factor for Ni(II) adsorption which estimates the nature of adsorption process with the determination of thermodynamic parameters. The thermodynamic parameter is an important aspect of Ni(II) adsorption study. To estimate spontaneity, thermal feasibility, and nature of the adsorption reaction, free energy change (ΔG0), enthalpy change (ΔH0) and entropy change (ΔS0) is evaluated using the following equations [95]: (1)

G 0 = H 0T S 0

ln(K c ) =

H0 1 + R T

S0 R

(2)

Where, Kc is the ratio between the amount of Ni(II) adsorbed and Ce indicates equilibrium concentration of Ni(II) and R represents the universal gas constant and T denotes the absolute temperature. ΔH0 and ΔS0 are considered from the slope and intercept of the plot of lnKC vc. 1/T respectively and then ΔG0 is estimated. Negative values of ΔG0 implies that Ni(II) adsorption process is spontaneous and thermodynamically favorable for the given experimental conditions. A positive value of ΔS0 indicates that the randomness of the solid-water interface has been increased and implies good affinity of Ni(II) towards adsorbents. A positive value of ΔH0 indicates that the reaction is endothermic [95]. This implies that as temperature increases, adsorption also increases. Table 8 displayed the various thermodynamic parameters for Ni(II) adsorption using various adsorbents [29,62,71,95,97,104,106,110,115,117,183,195]. From Table 8 it is noted that in most instances Ni(II) adsorption onto the substrate is spontaneous and endothermic. Entropies are typically positive values implying that Ni(II) adsorption primarily driven by a modification in structure or solvation rather than by a modification in chemical bonding.

9. Conclusion The study demonstrated the latest development in the potentiality of numerous adsorbents for Ni(II) adsorption from the aquatic environment. The macroscopic data, surface complexation model (SCM) and spectroscopic results of this investigation can provide a solid foundation for understanding surface-water behaviors and Ni(II) adsorption in the aquatic ecosystems. From the reviewed papers, it has found that very few articles have been published for competitive and comparative Ni(II) adsorption with other metal ions, anions, organic dyes and humic substances using various adsorbents. Thus, more research on Ni(II) adsorption is required. Reviewed papers also indicated that experimental conditions were typically taken into account during Ni(II) adsorption study and notably affected Ni(II) adsorption process. It is apparent in most reported papers that little or no Ni(II)

8. Future prospective research Although a significant number of research articles are publishing at the tremendous race on Ni(II) adsorption onto mineral surfaces, there are still numerous gaps which need more attention. The following facts can be considered for future research: 11

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Table 8 Calculation of thermodynamic parameters for Ni(II) adsorption onto various adsorbents at different temperature. Adsorbents

Temperature (K)

ΔG0 (kJ mol−1)

ΔH0 (kJ mol−1)

ΔS0 (kJ mol−1)

References

Manganese dioxide derived from ground water treatment sludge

298.15 308.15 318.15 292 303 313 333 293 303 313 333 296 313 333

−1.99 −2.57 −3.16 19.52 20.32 20.99 22.34 −1.34 −0.30 0.62 1.29 −4.28 −3.86 −3.73

15.49

58.63

[95]

4.07

67.09

[62]

−26.48

−6.31

[110]

11.58

24.65

[106]

303 313 323 300 308 313 318 323 298 313 333 298 308 318 298 308 318 298 308 318 303 308 318 323 333 338 293 323 343 303 318 333 303 318 333 298 313 323

−1.10 −1.26 −1.45 −1.08 −1.46 −1.82 −2.13 −2.27 −41.28 −50.18 −53.38 −5.78 −6.74 −7.32 −6.81 −6.83 −6.84 −6.84 −6.95 −7.09 −6.01 −7.60 −9.19 −10.78 −12.36 −17.13 −16.00 −18.11 −20.18 3.59 3.05 2.83 −8.43 −9.67 −10.48 −1.89 −1.72 −0.97

8.26

0.019

[29]

15.13

54.04

[104]

17.8

138.6

[115]

29.5

−77.84

[183]

18.45

−75.66

[183]

16.34

−73.45

[183]

90.22

0.318

[97]

8.43

0.83

[195]

11.33

25.70

[71]

12.48

69.24

[71]

15.3

43.91

[117]

Multi-walled carbon nanotubes (MWCTs)

Expanded perlite

Manganese oxide coated zeolite

Diatomite waste modified by EDTA Nigerian kaolinite clay

Polyaniline/CoFeC6N6 nanocomposite Unmodified modified hydrochar Acid-hydrochar Alkali-hydrochar Activated alumina

Calcium hydroxyapatite Rice husk ash (RHA) Carbon embedded silica from RHA L-methionine modified Dowex-50 ion exchanger

adsorption was noted at low pH while significantly higher Ni(II) adsorption was found at higher pH values depending on the surface charge of the substrates. Ni(II) formed an insoluble precipitate of Ni (OH)2 at pH 8.5 or higher. Impregnation of natural materials such as sand, clay and metal oxide into low-cost agricultural based substrates or nano-matrix showed excellent Ni(II) adsorption compared to their individual counterpart. Ni(II) adsorption data of isotherm, kinetics, and thermodynamics onto various adsorbents was reviewed and it resolved that Langmuir and Freundlich equations were frequently used to model adsorption isotherm data while PS1 and PS2 kinetic were also modeled commonly. Ni(II) adsorption data were best defined by the Langmuir model and PS2 in most cases. Ni(II) adsorption process onto various adsorbents was commonly endothermic and spontaneous in many cases. The Ni(II) adsorbed onto substrates typically by forming outersphere Ni(II)-complexes at low pH and by forming inner-sphere Ni(II) complexes and precipitate formation at higher pH values. The Ni(II)

adsorption involved more than one surface reaction. Commonly Ni(II) adsorption was found to decrease in presence of another metal ion, anion, dye and humic acid(HA) due to the competition of both species for a finite number of surface sites or development of soluble Ni(II)-HA complexation in solution. Adsorption of Ni(II) onto substrates has an important influence on their bioavailability, transport, and toxicity in ecological conditions. However, the reported experimental conditions in most previous papers are slightly dissimilar from real effluent. The results in this investigation may provide essential assumptions for identifying the change tendency on Ni(II) in real samples and have a significant application in removing metal ions. Declaration of Competing Interest The authors declare that they have no known competing financial 12

Journal of Environmental Chemical Engineering 7 (2019) 103305

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interests or personal relationships that could have appeared to influence the work reported in this paper.

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