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Journal of Hazardous Materials 248– 249 (2013) 1– 19

Contents lists available at SciVerse ScienceDirect

Journal of Hazardous Materials

jou rn al h om epage: www.elsev ier .com/ loc ate / jhazmat

eview

efluoridation of drinking water using adsorption processes

aripurnanda Loganathana, Saravanamuthu Vigneswarana,∗, Jaya Kandasamya, Ravi Naidub

Faculty of Engineering and Information Technology, University of Technology, Sydney, NSW, 2007, AustraliaCentre for Cooperative Research for Contamination Assessment and Remediation of the Environment (CRC CARE), University of South Australia, Adelaide, SA 5095, Australia

i g h l i g h t s

Comprehensive and critical literature review on various adsorbents used for defluoridation.pH, temperature, kinetics and co-existing anions effects on F adsorption.Choice of adsorbents for various circumstances.Adsorption thermodynamics and mechanisms.Future research on efficient, low cost adsorbents which are easily regenerated.

r t i c l e i n f o

rticle history:eceived 2 August 2012eceived in revised form8 December 2012ccepted 26 December 2012vailable online 4 January 2013

eywords:

a b s t r a c t

Excessive intake of fluoride (F), mainly through drinking water, is a serious health hazard affectinghumans worldwide. There are several methods used for the defluoridation of drinking water, of whichadsorption processes are generally considered attractive because of their effectiveness, convenience,ease of operation, simplicity of design, and for economic and environmental reasons. In this paper, wepresent a comprehensive and a critical literature review on various adsorbents used for defluoridation,their relative effectiveness, mechanisms and thermodynamics of adsorption, and suggestions are madeon choice of adsorbents for various circumstances. Effects of pH, temperature, kinetics and co-existing

dsorptionefluoridationluorideluoride toxicityater treatment

anions on F adsorption are also reviewed. Because the adsorption is very weak in extremely low or highpHs, depending on the adsorbent, acids or alkalis are used to desorb F and regenerate the adsorbents.However, adsorption capacity generally decreases with repeated use of the regenerated adsorbent. Futureresearch needs to explore highly efficient, low cost adsorbents that can be easily regenerated for reuseover several cycles of operations without significant loss of adsorptive capacity and which have goodhydraulic conductivity to prevent filter clogging during the fixed-bed treatment process.

© 2013 Elsevier B.V. All rights reserved.

ontents

1. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 22. Adsorption mechanisms . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 23. Factors influencing adsorption. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4

3.1. pH . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 43.2. Co-existing anions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 43.3. Temperature . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 113.4. Adsorption kinetics . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11

4. Adsorbents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 114.1. Metal oxides and hydroxides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 114.2. Layered double hydroxides. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12

4.3. Ion exchange resins and fibres . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

4.4. Zeolites . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .4.5. Carbon materials . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

∗ Corresponding author. Tel.: +612 9514 2641.E-mail address: s.vigneswaran@uts.edu.au (S. Vigneswaran).

304-3894/$ – see front matter © 2013 Elsevier B.V. All rights reserved.ttp://dx.doi.org/10.1016/j.jhazmat.2012.12.043

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14

2 P. Loganathan et al. / Journal of Hazardous Materials 248– 249 (2013) 1– 19

4.6. Natural materials . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 144.7. Industrial by-products. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15

5. Adsorption thermodynamics . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 166. Fluoride desorption and adsorbent regeneration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 177. Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18

Acknowledgements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18 . . . . . .

1

itacciadctWilStw

acrotbt

uinftaf

iNetTaFtr

ihsrFbit

t

References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

. Introduction

Fluoride (F) has beneficial effects on teeth at low concentrationsn drinking water (0.4–1.0 mg/L), especially for young children inhat it promotes calcification of dental enamel and protects teethgainst tooth decay. Excessive levels of F on the other hand canause a number of problems ranging from mild dental fluorosis torippling skeletal fluorosis as the level and period of exposure to Fncreases [1]. The World Health Organisation [2] has recommended

guideline value of 1.5 mg/L as the concentration above whichental fluorosis is likely. However, it is also important to considerlimatic conditions and quantity of water intake and other fac-ors such as F intake from certain diets when establishing F limits.

ater consumption in hot humid regions is generally higher thann temperate regions and therefore the F concentration limit forikely fluorosis should be lower. For example, the US Public Healthervice [3] has recommended that the upper limit for F concen-ration in drinking water should be decreased from 1.7 to 0.8 mg/Lith increases of the average maximum daily air temperature.

High F intake has been suspected being involved in a range ofdverse health problems in addition to fluorosis, including can-er, impaired kidney function, digestive and nervous disorders,educed immunity, Alzheimer’s disease, nausea, adverse pregnancyutcomes, respiratory problems, lesions of the endocrine glands,hyroid, liver and other organs [1,4–9]. However, there appears toe no convincing evidence for F being directly involved in causinghese conditions [1].

Elevated concentration of F in drinking water is due to its nat-ral occurrence or industrial activities. Many rocks and minerals

n the earth’s crust contain F [10,11] which can be leached out byatural weathering and rainwater, causing F contamination of sur-

ace and ground waters. Besides this natural source, F also entershe water bodies from waste waters produced from industries suchs aluminium, steel, glass, semiconductors, electronic, tooth paste,ertiliser and insecticide manufacturing plants [9,12–15].

Fluorosis due to excessive concentration of F has been reportedn at least 28 countries from South Asia, Africa, the Middle East,orth, Central and South America, and Europe [1]. In India, it wasstimated that 56.2 million people were affected by fluorosis andhis problem was prevalent in 17–19 out of the 32 States [16,17].he major source of F in a majority of countries is rocks and miner-ls such as fluorospar, cyyolite, and fluorapatite containing F [17].or example, Choi and Chen [18] reported extremely high F concen-ration (>1000 mg/L) in surface water in areas with F-rich volcanicocks.

One method of reducing excessive concentrations of F in waters to blend water having a high F concentration with water thatas a low F concentration from an alternative source. If such aource is not readily available, defluoridation is the only meansemaining to prevent fluorosis [1]. For contaminants other than, water treatment methods are used to remove contaminants toelow the maximum level permissible but defluoridation is special

n that the treated water should have an optimum F concentrationo derive the beneficial effects of F.

The main methods of defluoridation are precipita-ion/coagulation, adsorption, ion exchange, reverse osmosis, and

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18

electrodialysis. Of these, precipitation/coagulation and adsorp-tion are convenient methods and are widely used, especially indeveloping countries’ rural areas. The scale and treatment sitediffer between industrialised countries and developing countries.In industrialised countries the treatment of water is generallyperformed at water treatment plants close to the water source butin developing countries it is carried out at a village communitylevel or at a household level [1] using simple inexpensive locallyavailable adsorptive media [9,19–21]. Industrialised countriesgenerally use more efficient but more costly adsorption mediaincluding synthetic ion exchange resins as well as advancedtechniques such as reverse osmosis and electrodialysis.

The advantages and shortcomings of the various methods ofdefluoridation are presented in Fig. 1. Of these various methods thatof adsorption is generally considered attractive because of its effec-tiveness, convenience, ease of operation, simplicity of design, andeconomic and environmental considerations, provided low-costadsorbents which can effectively remove F around the neutral pHof drinking water are used. The precipitation/coagulation methodwhere lime and Al salts are used to remove F as a CaF2 precipitatefollowed by removal of left over F in solution by co-precipitationwith and adsorption on to the precipitated Al(OH)3 is further devel-oped into a technique called Nalgonda process in India [17,24,25].This process is extensively used in India and Africa [1]. However,the main drawback of this technique is the low effectiveness of Fremoval and production of toxic AlF complexes in solution [8].

Previous reviews on defluoridation of water presented timelyinformation on several adsorbents used for defluoridation, but theygenerally did not focus on the chemical mechanisms and solutionfactors influencing the adsorption processes, and regeneration ofthe adsorbents [1,17,26,27]. The objective of this paper is to com-pile and present current information on the potential of the variousadsorbents used for defluoridation of drinking water, their relativeeffectiveness, mechanisms and thermodynamics of adsorption, fac-tors influencing adsorption and methods of adsorbent regenerationfor reuse. Based on the review adsorbents are selected for specificcircumstances.

2. Adsorption mechanisms

The capacity, energy and kinetics of adsorption of F byadsorbents are controlled by the mechanism of adsorption. Under-standing these mechanisms can provide useful information onthe optimisation of the adsorption process in water treatmentplants and the subsequent desorption/adsorbent regeneration pro-cess for reuse. There are mainly five mechanisms of F adsorption,namely: (1) van der Waals forces (outer-sphere surface complex-ation, (2) ion exchange (outer-sphere surface complexation), (3)hydrogen bonding (H-bonding) (inner-sphere surface complexa-tion), (4) ligand exchange (inner-sphere surface complexation),and (5) chemical modification of the adsorbent surface. The firsttwo mechanisms are governed by weak physical adsorption and

are non-specific to F, whereas the third and fourth mechanismsare governed by strong chemical adsorption specific to F. The fifthmechanism is governed by both specific and non-specific adsorp-tion. In the presence of the other anions in water, F cannot be easily

P. Loganathan et al. / Journal of Hazardous Materials 248– 249 (2013) 1– 19 3

Prec ipitatio n/coagu lation

Adsorption /Ion Exchange

Most widely use d

Medium cost

Low effec tiveness ; cann ot

remove F below 5 mg/L

beca use of high so lubi lity

product of C aF2; need

secondary t rea tment

Lar ge amounts of chemicals

required

Precise control of chemicals

addi tions ( frequ ent test ing of

feed and treated water)

Cost s of chemicals, ch emic al

storage and feedi ng syst em

Lar ge volu mes of waste slud ge;

disposal problem

Acid neutra lisation o f t reated

wate r requir ed

Toxic chemicals left in treated

wate r (Al F complexes, SO4)

Most wid ely use d

Can b e cost ly (esp ecially ion -

exchang e resin s), but can use

low-cost adso rbent (in clud ing

certain waste m ate rials)

Effective even at low F

concent ration

Simpl icity and f lexib ility of

desi gn

Ease of o peration

No waste p rod uction

Low sele ctivi ty against some

anions for adsorption /all anions

for ion e xchange com peting ions

Frequent adsorbent regeneration

or repl acement requir ed

Granula r adsorb ent better for

good hydraul ic flow

Effective , mos tly at pH < 7 fo r

adsorption

Rever se osmosis

Ele ctrodia lysis

Exce llent r emoval

Very h igh capi tal cost. Very

high op eration al ( energy) cost

No chemica ls requir ed

No waste p rod uction

No ion selectively; benefi cial

nutrients and othe r cont aminants

removed togeth er with F

Some membr anes pH s ensi tive

F con centr ated residue disposal

problem

Wate r w asted

Cloggin g, sca ling and f oul ing

problems

Exce llent removal

High capi tal cost. H igh

operation al (energy) cost

No che micals required

No waste p rod uction

No ion s ele ctively; benef icial

nutri ents and othe r cont aminants

removed t ogeth er with F

Ski lled l abour required

Polarization problem

fluori

ratoa

taaswmaf

lbore

Fig. 1. Common technologies for de

emoved by adsorbents using the first two mechanisms. Fluoridedsorption resulting from the third and fourth mechanisms selec-ively removes F from water in the presence of most types of anions;nly anions (e.g. phosphate) which also adsorb specifically on thedsorbent compete with F for adsorption sites.

Van der Waals forces are weak short range forces acting betweenwo atoms. The larger the adsorbate size the greater the force ofttraction. Therefore adsorbates with high molecular weights suchs dissolved organic matter are adsorbed on adsorbents having highurface area through van der Waals forces. This is the reason for theeak adsorption of F and strong adsorption of dissolved organicatter on activated carbon [28]. Fluoride was considered to be

dsorbed on manganese oxide-coated alumina by van der Waalsorces at high pH [29].

Ion exchange is a stoichiometric process where any counter ioneaving the ion exchanger surface is replaced by an equivalent num-

er of moles of another counter ion to maintain electro-neutralityf the ion exchanger. The ions are adsorbed physically by fullyetaining their inner hydration shell and the adsorption is due tolectrostatic or Coulombic attraction. The ion exchange process is

dation of drinking water [8,22,23].

rapid and reversible. Fluoride removal by ion exchange resins [30]and ion exchange fibres [31] is mainly governed by the ion exchangemechanism as illustrated in Fig. 2a. Ion exchange tends to pre-fer counter ions of higher valency, higher concentration and ionsof smaller hydrated equivalent volume [32]. Therefore F removalusing ion exchange resins is difficult because the order of selectiv-ity for anions by anion exchange resins is as follows: citrate > SO4

2−,oxalate > I− > NO3

− > CrO42− > Br− > SCN− > Cl− > acetate > F− [32].

H-bonding is a strong dipole–dipole attractive force betweenbonding of the strong electropositive H atom in a molecule in anadsorbent or adsorbate and a strong electronegative atom such asoxygen or fluorine in another molecule [33]. The energy of adsorp-tion in H-bonding is stronger than in van der Waals forces andion exchange but weaker than in the ligand exchange process dis-cussed in the next paragraph. H-bonding occurs in the adsorptionof F on ion exchange resins [30,34] and coal-based adsorbents [35]

as illustrated in Fig. 2b.

In a ligand exchange mechanism, the adsorbing anion suchas F− forms a strong covalent chemical bond with a metalliccation at the adsorbent surface resulting in the release of other

4 P. Loganathan et al. / Journal of Hazardous

F

ptcapooAatostl

ctsIAtbcc

ahesdIoo

The extent of the competition depends on the relative concentra-

ig. 2. Mechanisms of F adsorption (�, adsorbent; Me, multivalent metallic cation).

otential determining ions such as OH− ions previously bonded tohe metallic cation (Fig. 2c). Thus, F is said to form an inner sphereomplex or is specifically adsorbed on the adsorbent surface. Thedsorption of F on several multivalent metal oxides near neutralH was reported to have increased the pH of solutions as a resultf release of OH− ions from the adsorbents by ligand exchangef OH groups on the adsorbent surface with F in solution [36,37].dsorbents with a ligand exchange mechanism have the particulardvantage of combining high adsorption capacity with high selec-ivity for anions. These adsorbents can remove large proportionsf anions having higher selectivity for adsorption from very diluteolutions of the anions even in the presence of higher concentra-ions of competing anions of lower selectivity. Adsorption of F byigand exchange is illustrated in Fig. 2c [23,37,38].

The adsorption capacity of F on adsorbents can be increased byhemical modification of adsorbent surfaces (Fig. 2d). This is par-icularly of advantage in the case of adsorbents possessing negativeurface charges which tend to repel the similarly charged F− ions.n such adsorbents, positively charged multivalent cations such asl3+, La4+, Zr4+, Fe3+, and Ce3+ are impregnated onto the adsorbent

o create positive charges on the adsorbent surface for attracting F−

y coulombic forces as well as producing adsorption sites capable ofhemical interaction with F [13,23,39–44] (Fig. 2d). These metallications act as a bridge in adsorbing F onto the adsorbent.

Adsorption of F on many microporous adsorbents is recogniseds a two-step process; an initial rapid adsorption (mostly within anour) at the outer surface of the adsorbent that reaches a pseudo-quilibrium at the solid-solution interface followed by a muchlower process (hours to days) where the F moves by intra-particleiffusion into the interior pores and channels of the adsorbent [45].

ntra-particle diffusion rate is directly related to the square rootf time of adsorption. Therefore if a straight line relationship isbtained between the rate of adsorption and square root of time

Materials 248– 249 (2013) 1– 19

with the line passing through the origin, it can be inferred thatthe diffusion process controls the adsorption, especially at longertimes. Such a relationship was obtained in many studies for theadsorption of F on different adsorbents (granular ferric hydroxide[7], manganese oxide-coated alumina [29], fly ash [46], activatedalumina [47], and alum sludge [48]).

It could be concluded that ligand exchange is the predomi-nant mechanism of F adsorption for inorganic adsorbents havinghigh adsorption capacities. For organic adsorbents having highadsorption capacities, H-bonding seems to be the predominantmechanism.

3. Factors influencing adsorption

3.1. pH

As for other contaminants, the removal of F from water by adsor-bents is influenced by several factors including pH, co-existing ions,temperature, adsorption kinetics, and adsorbent particle size andactivation. Of these, pH is generally considered to be the mostimportant factor. Fluoride adsorption is low at both very low andvery high pH. The pH at which the maximum amount of F isremoved depends on the adsorbent characteristics but generallyit is between 4 and 8 (Table 1). One of the important propertiesof the adsorbent influencing the extent of F adsorption is the pHat the point of zero charge (PZC). For example, Fe and Al oxideshaving a PZC at around 7–8 remove the maximum amount of Fat pH 6–8 [7,47,49] and activated carbon with a PZC of 3.9–4.7 wasreported to have removed the maximum amount of F at pH 3–4 [42].At pH values above the PZC the surface of adsorbents is negativelycharged, therefore the negatively charged F ions are not attracted tothe adsorbent surface. At pH values lower than the PZC the surfaceis positively charged, therefore F ions are adsorbed. In some situa-tions, at low pH values, F exist as positively charged AlF complexesand this can reduce F adsorption [11,49,75].

The main reason for a reduction in F adsorption below pH 4 isthat F forms weakly ionised HF at these pH values [13,47,48,50].At a pH above 7–8 the removal of F decreases not only because theadsorbent surface becomes negatively charged but also because theconcentrations of hydroxyl, bicarbonate and silicates ions increaseso that they compete with F for adsorption. In addition to increasednumber of positive surface charges at low pH values increasing Fadsorption, surface protonation at a pH less than the pH of PZCprovides an increased number of H atoms at the adsorbent surfaceleading to an increased number of H-bonding between the H atomsand F in solution resulting in increased F adsorption [76]. Adsorp-tion by a ligand exchange mechanism is also favoured at low pHbecause of a stronger attractive force between F and the adsorbentsurface and the presence of more hydroxylated sites for exchangewith F than at a high pH [4,29,61].

It is apparent from the literature review that F adsorption is gen-erally lowest at extremely low and high pH values. The adsorptionis highest around the neutral pH that is commonly found in naturalwater. Therefore prior pH adjustment is not normally required foreffective removal of F in treatment plants.

3.2. Co-existing anions

In natural water, several anions including PO43−, Cl−, SO4

2−,Br−, NO3

− are simultaneously present with F at different concen-trations which can compete with F for removal by adsorbents.

tions of the ions and their affinity for the adsorbent. Meenakshiand Viswanathan [30] reported that increased concentrations ofCl−, SO4

2−, Br−, NO3−, and HCO3

− decreased F adsorption by an

P. Loganathan

et al.

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azardous M

aterials 248– 249 (2013) 1– 19

5

Table 1Characteristics of adsorptive removal of fluoride from water by different adsorbents.

Adsorbent Adsorptionmethod: batch (B),column (C); Watertype: synthetic (S),wastewater (W)

pH; temperature(◦C)

Initial (I), equilibrium (E)concentration (mg/L); Adsorbentconcentration (g/L); column ht,height; d, diameter; FL, flow rate

Adsorption (Ads) capacity (mg/g) andother results. Maximum (max)

Best kinetic modelto fit data;Equilibrium (equil)

Best equilibriummodel to fit data

Reference

Metal oxides and hydroxidesGranular ferrichydroxide

B; 3–12; I, 1–100; Max ads pH 3–8, then decreased withincreased pH to 12; Langmuir ads max(pH 6–7) 10 ◦C, 3.68; 25 ◦C, 5.97

Pseudo-first order,Bangham model

Langmuir [7]

S 10, 25 2

Manganese oxidecoated alumina

B, C; B, 3–12; B, I, 6.2–42.1; B, Max ads pH 4–6, decreased withincreased pH from 6 to 12. Langmuirads max 7.1 at pH 5.2

Pseudo-first,pseudo-second,diffusion

Langmuir [29]

S 25 C, I 5; 140; 50 cmht, 2.4 cmd, FL2.39 m3/m2 h

C, breakthrough point 669 bed volume

C, 5.2;25

Manganese oxidecoated alumina

B, C; B, 3–12; B, I, 2.5–30 B, Max ads pH 4–7, lowest at pH 12;Langmuir ads max 2.85 at pH 7

Pseudo-first,pseudo-second,diffusion

Langmuir [36]

S 30 C, 3.56;C, 7; 0.5 × 0.028 md, 0.3 mht, FL

2.19 m3/m2 hC, Bed saturation F concentration 1.25g/L

30

Activated alumina B, C; B, 3–12; B, I, 2.5–30 B, Max ads pH 4–7, lowest at pH 12;Langmuir ads max 1.08 at pH 7

Pseudo-first,pseudo-second,diffusion

Langmuir [36]

S 30C, 7; C, 3.56; C, Bed saturation F concentration

0.47 g/L30 0.5 × 0.028 md, 0.3 mht, FL

2.19 m3/m2h

Activated alumina B, C; B, 4–10; B, I, 2.5–14; Max ads pH 7 Surface adsorption,diffusion

Langmuir andFreundlich

[47]

S Room temp. 4–40 Langmuir ads max 2.41 at pH 7C, 7; room temp. C, 5; 200

550 mmht, 51 mmd, FL20,30 mL/min

96% removal

Activated alumina B; 4–11; – Max ads pH 5–7 Pseudo-secondorder

Langmuir andFreundlich

[49]

S 30 12.5 Langmuir ads max 16.3 at pH 6

Rare earth oxides B; 3–11; E 1–30; Max ads pH 6–6.5 First 5 min most Fads; 40 min equil

Langmuir [50]

S 29 1–8 Langmuir ads max 196 at pH 6.5

Hydroxyapat-ite (HA) B; – I, 19–19,000; At initial F concentration of ≤190 mg/L,60–80% F removed by porous HA and30–35% by crystalline HA

– – [51]

S Room temp. 10

6P.

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Journal of

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Materials

248– 249 (2013) 1– 19

Table 1 (Continued)

Adsorbent Adsorptionmethod: batch (B),column (C); Watertype: synthetic (S),wastewater (W)

pH; temperature(◦C)

Initial (I), equilibrium (E)concentration (mg/L); Adsorbentconcentration (g/L); column ht,height; d, diameter; FL, flow rate

Adsorption (Ads) capacity (mg/g) andother results. Maximum (max)

Best kinetic modelto fit data;Equilibrium (equil)

Best equilibriummodel to fit data

Reference

Hydroxyapat-ite (HA) B; 6; I;2.5 × 10−5–6 × 10−2; >90% F adsorption; Langmuir ads max4.54 compared to fluorspar 1.79,activated quartz 1.16, calcite 0.39,quartz 0.19

Pseudo-secondorder

Langmuir [52]

S Room temp. 17

La impregnated silicagel

B; B 4–9; B, I, 10.4–104 Ads max pH 4.7–7.3, Langmuir ads max3.8 at pH 7

Pseudo-firast order Langmuir [43]

W 20 0.1C; C 7.5; C, I, 72; 4; 9 cm ht, 1 cm d, FL

0.5 mL/minColumn ads capacity 3.4

W 20

Layered double hydroxides (LDH)Mg/Al LDH calcinedand uncalcined

B; 7; E, 0–70; Calcination increased F ads (130 ◦Coptimum). Ads capacity 35 atequilibrium F concentration 70 mg/L

Adsorption reachedmax at 15 min

– [53]

S 25 2.5

Mg/Al LDH calcined(200–800 ◦C)

B; 5–10; E, 0–250; Max ads pH 6.0; Langmuir ads max 213at pH 6. Ads highest for calcinationtemp. 500 ◦C. Ads Mg/Al LDH >Ni/Al,Zn/Al LDHs, Mg/Al molar ratio 2:1 best

– Freundlich [54]

S 30 1

Zn/Al LDH calcined(450 ◦C)

B; 4–10; E, 2–60; Max ads pH 6; Langmuir ads max 17 atpH 6. Max ads 13.43 for adsorbent dose0.2 g/L, F concentration 10 mg/L, pH 6

ads reached equilin 4 h

Langmuir [55]

S 30 1

Na/Mg/Al LDH calcined(500 ◦C)

B; S: 5,7,9; S: I, 5; Distribution coefficient between solidand solution, S: pH 5, 3501; pH 7,1047; pH 9, 1155

80–97% F removalin 1 h

– [56]

S, well water Room temp. 10 W: 414–437W:8.4,8.5 W: I, 5.9, 6.9;Room temp. 10

Ion exchange resins and fibresThio-urea modifiedAmberlite (TUA)

B, C; B: 1–10; B: E, 2–13; B: ads max at pH 7 for TUA and pH9–10 for amberlite (A). At pH 7, adscapacity of TUA three times that of A.Langmuir max for TUA at pH 7, 61.

Ads reached equilin 30 min

Langmuir [34]

S 25 10 C: column ads capacity 50C: 7; C: I, 16;25 2; 0.8 mmd, 50 mmht, FL,1 mL/min

Chelating resin (CR),anion exchange resin(AER)

B; 3–11; I, 2–10; pH had no significant effect on ads;Langmuir ads max at pH 7: CR 1.3, AER1.5. At low F concentration CR removed30% more F than AER. Field water: CRhad higher F removal

Pseudo-secondorder

Langmuir [30]

S, field water 30 20

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et al.

/ Journal

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aterials 248– 249 (2013) 1– 19

7

Table 1 (Continued)

Adsorbent Adsorptionmethod: batch (B),column (C); Watertype: synthetic (S),wastewater (W)

pH; temperature(◦C)

Initial (I), equilibrium (E)concentration (mg/L); Adsorbentconcentration (g/L); column ht,height; d, diameter; FL, flow rate

Adsorption (Ads) capacity (mg/g) andother results. Maximum (max)

Best kinetic modelto fit data;Equilibrium (equil)

Best equilibriummodel to fit data

Reference

Metals loadedAmberlite resin

B, C; B: 1–8; B: E, 0–60; B: Langmuir ads max at pH 7, resinwith La, Ce, Al, 25; Fe, 49; Y, 19. pH forads max: Ce 4–7, Fe 3, Al 5–9, L 3–7

Ads equil time: La2 h; Ce 1 h; Fe, Al, Y>5 h

Langmuir [39]

Field water 30 1.6 C: column ads capacity for La resin atpH 6, 20. Bed saturation 300 bed vol.

C: 6.0; C: I, 7.9;30 2; 8 mmd, FL 6 mL/h

Al-Amberlite resin B, C; B: 4–9.1; – B: max ads at pH 4–7, low ads pH 9,Langmuir ads max 4.6 at pH 4

Pseudo-secondorder

Langmuir [57]

S 30 – C: column ads capacity at pH 5.5 and6.7 were 1.13 and 0.72 (FL 460 mL/h)

C: 5.5, 6.7 C: I, 40;30 20–50;

2 cm d, 16 cmht, FL 280–700 mL/h

Ion exchange fibre B, C; B: 2–6; B: E, 0–60; B: ads max at pH 3.0 and decreasedwith increased pH to 6; Freundlichadsorption constantPO4

3− > AsO43− > F−

Pseudo-secondorder

Freundlich [31]

S 25 4 C: column ads capacity for PO43,

AsO43− , F− 156, 96, 45, respectively

C: –; C: I, 10; Equil reached in5 min

25 11.5 cmd, 5 cmht, FL 2.5–3.5 mL/min

Al-chelating porousanion exchanger

B, Tap water 2–9; E, 0–100; Ads max at pH 2.6–6.9. Ads capacity 30for 5 g adsorbent at E = 100 mg/L

1 h equil time (95%adsorption)

– [58]

30 5

ZeolitesAl, La, Zr loaded naturalzeolites (Z)

B; – S: I, 1–20; S: Langmuir ads max: Zr (Z) 3.4–4.1,La(Z) 2.4–2.6, Al(Z) 2.0–2.4

– Langmuir [23]

S, tap water (T) 30 2 T: F removal(%): Zr(Z) 91.1, La(Z) 90.4,Al(Z) 89.7

T: I, 2.9;6

Al loaded foursynthetic zeolites

B; 2–11; I, 5–80; Ads max pH 4–8; F ads at equilibriumconcentration 40 mg F/L and pH 4–6was 6–16

Elovich Redlich-Peterson [41]

S 25 0.05–0.2

Al, La loaded syntheticzeolites

B; S: 3.5–9; S: I, 10–80; S: ads max at pH 6–9; Langmuir adsmax 20 ◦C, Al(Z) 34, La(Z) 45.At 40 mg/Lequilibrium conc., ads capacity 16 vs 8for Al(Z) vs La(Z). (F): F conc reductionmore by Al(Z) than La(Z)

- Redlich-Peterson [40]

S, field water (F) 20–40 2F: 7.4; F: I, 3.3, 4;30 1–4

8P.

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Materials

248– 249 (2013) 1– 19Table 1 (Continued)

Adsorbent Adsorptionmethod: batch (B),column (C); Watertype: synthetic (S),wastewater (W)

pH; temperature(◦C)

Initial (I), equilibrium (E)concentration (mg/L); Adsorbentconcentration (g/L); column ht,height; d, diameter; FL, flow rate

Adsorption (Ads) capacity (mg/g) andother results. Maximum (max)

Best kinetic modelto fit data;Equilibrium (equil)

Best equilibriummodel to fit data

Reference

Carbon materialsKMnO4 modifiedactivated carbon

B; 2–10; E, 5–20; Ads max at pH 2, then decreased withpH. Langmuir ads max at 25 ◦C, 15.9.Ads decreased with temperatureincrease

Pseudo-secondorder

Langmuir–Freundlich

[59]

S 25, 45, 55 –

Bone char, activatedcarbon, carbon black

B; 4.6–9.2; E, 0–20 Ads capacity of bone char increasedwith pH. Freundlich adsorptionconstant highest for bone char andlowest for carbon black

– Freundlich [60]

S 25

Aligned carbonnanotubes (ACNT)

B; 3–11; E, 0–16 Ads max at pH 7; ads capacity atequilibrium concentration 10 mg F/Lfor activated carbon, � Al2O3, andACNT: 0.32, 3.7, and 4.1, respectively

– Freundlich [12]

S 25

Alumina loaded CNT(calcined, 250–1050 ◦C)

B; 3–11; – Ads max at pH 6–9; ads max of 9.6 atpH 6 and calcined temp. 450 ◦C for0.5 mg/L adsorbent and initialconcentration 6 mg F/L

Equil reached at20 h

Freundlich [61]

S 25

Wood, animal, fishbone activated charcoal

B; 2–12; I, 4–30; Ads decreased from pH 2 to pH 12; Fishbone charcoal had the highest F ads

Most of F removedat 2 h

Langmuir,Freundlich

[62]

S – –

Al impregnatedactivated carbon

B; 3–7; I, 0.5–15 Ads decreased from pH 3 to 7.Calcining at 300 ◦C gave the highestads among 300–1000 ◦C. Langmuir adsmax 1.07, plain carbon 0.49

– Langmuir,Freundlich

[42]

S 25 E, 0–6.5;–

Natural materialsChitosan/ B; 3–11; I, 9–15; Ads max at pH 3, decreased to pH 11.

At pH 7, 10 mg F/L, 2 g adsorbent/L,composite ads 1.3; LDH alone 1.0,chitosan alone 0.05; Langmuir ads max1.9 for composite

Pseudo-secondorder, diffusion

Langmuir [63]

LDH composite S 30 2Chitin/nano-hydroxyapatitecomposite

B; 3–11; I, 6–12; Ads max at pH 3, decreased withincreased pH to pH 11; Langmuir adsmax 8.4 at pH 7

Pseudo-secondorder, diffusion

Langmuir [64]

S 30 2 Freundlich

Chitosan/nanohydroxyapatitecomposite

B; 3–11; I, 9–15; Ads max at pH 3, decreased withincreased pH to pH 11; At pH 7, 10 mgF/L, adsorbent 5 mg/L, composite adscapacity 1.56, hydroxyapatite 1.30,chitosan 0.05; Langmuir ads max 2.04

Pseudo-secondorder, diffusion

Langmuir [65]

S 30 2

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/ Journal

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aterials 248– 249 (2013) 1– 19

9

Table 1 (Continued)

Adsorbent Adsorptionmethod: batch (B),column (C); Watertype: synthetic (S),wastewater (W)

pH; temperature(◦C)

Initial (I), equilibrium (E)concentration (mg/L); Adsorbentconcentration (g/L); column ht,height; d, diameter; FL, flow rate

Adsorption (Ads) capacity (mg/g) andother results. Maximum (max)

Best kinetic modelto fit data;Equilibrium (equil)

Best equilibriummodel to fit data

Reference

Magnetic (Fe)-chitosan B; S: 5–9; I, 5–140; pH no significant effect. Ads higherthan activated alumina. Langmuirtwo-site max ads 24

Equil at 90 min;Pseudo-secondorder

Langmuir (oneand two sites)

[66]

S 30 1 Bradley

La loaded chitosan B; S: 5–9; S: E, 0–15; S: Ads max pH 6.7; max ads 5.5 at pH6.7 and equilibrium conc.15 mg F/L

Pseudo-first order,diffusion

– [67]

S, field water (F) 30 – F: ads capacity 1 compared to 2 fordistilled water at equilibrium conc.8 mg F/L

F: 7; F: I, 10.2;30 –

Natural chitosan B; 2–10; – Max ads at pH 6; Langmuir ads max1.39 at pH 6

Equil at 5 min Langmuir, [22]

S 20 Freundlich

Zr loaded collogen B; 3.5–11; I, 19–95; Max ads pH 5–8, drastic decrease frompH 9 onwards. Langmuir ads max 2.18at pH 5–8

Equil at 500 min Langmuir [13]

S 30 1

La loaded cross-linkedgelatin

B; 2–12; E: 0–50; pH 2–5, >90% ads, then decreased withincreased pH to pH 12 (17%); Langmuirads max 21.3 at pH 5–7

Equil at 40 min;pseudo-first order

Langmuir [44]

S 29 4

Laterites with Ni B; S: 2.5–10; S: E, 5–40; S: ads max at pH 5–6 and 3–5;Langmuir ads max 12–15 at pH 5

Pseudo-first order, Langmuir [68]

S, field water 32 2 F: 3 repeated stages of ads reduced Fconcentration to <0.5 mg/L

Diffusion

F: 7.75; F: I, 10.25;32 14

Al, Fe, Ca loaded siltyclay (C)

B; Al, Ca-C, C alone:7–8; Fe-C 2.7–3.0;30

I, 10; % F removal: C alone 25, Al-C 78–94Ca-C 32–52, Fe-C 94–98

– – [69]

S 100

Soil with high Fe oxide(calcined 500 ◦C)

B, C; B: 7; B:I, 4.8–95; B: max ads 1 at equilibriumconcentration 60 mg/L

– Freundlich [38]

S 25 33C: 7; C: I, 4; C: F concentration reached 1 mg/L after

120 pore volumes. Column adscapacity 0.15

25 47; 50 mL column, FL 0.2 mL/minVolcanic ash soil B; – I, 0–50; At equilibrium concentration of

19 mg/L, ads capacity was 2.9Equil in ∼2 h at lowconc, >24 h at highconc

Langmuir [21]

S 18 40 Langmuir ads max 5.5

Bentonite (Be),kaolinite (Ka)

B; 2.5–9.5; I, 5; Increased pH, decreased ads, Be:46–29%, Ka: 38–5% for 2 g/L adsorbent

– – [70]

S 32 1–8

10P.

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Journal of

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Materials

248– 249 (2013) 1– 19

Table 1 (Continued)

Adsorbent Adsorptionmethod: batch (B),column (C); Watertype: synthetic (S),wastewater (W)

pH; temperature(◦C)

Initial (I), equilibrium (E)concentration (mg/L); Adsorbentconcentration (g/L); column ht,height; d, diameter; FL, flow rate

Adsorption (Ads) capacity (mg/g) andother results. Maximum (max)

Best kinetic modelto fit data;Equilibrium (equil)

Best equilibriummodel to fit data

Reference

Industrial by-productsBasic oxygen furnaceslag (BOFS)heated1000 ◦C

B; 2–10; I, 1–50; Ads increased from 21 to 93% as pHincreased from 2 to 7, then nearlyconstant; Langmuir ads max at 25 ◦C,4.6 and 45 ◦C, 8.1

Equil at 35 min;pseudo-first order

Langmuir [71]

S 25 5

Granular red mud B; B:2.5–7.3; B: I, 5–150; B: ads max at pH 4.7; Langmuir adsmax 8.92 at pH 4.7

Equil at 6 h;pseudo-secondorder

Freundlich,Redlich-Peterson

[72]

S 25 2.5 C: column total ads capacity 2.05 (0.64by batch trial for initial F, 5 mg/L

C; C: 4.7; C: I, 5;10, 0.635 cm2 area,15 cmht,FL 2 mL/min

S 25

Activated red mud (A),unactivated (UA)

B; 1–10; I, 100–1000; Ads max at pH 5.5; ads capacity A 4.8,UA 1.0 at equilibrium concentration20 mg/L; Langmuir ads max A 6.3, UA3.1 at pH 5.5

Equil at 2 h Langmuir [4]

S – 1–8.4

Waste carbon slurry450 ◦C (activated)

B; B:2–11; I, 1–11;1 B: ads max at pH 7.58; Langmuir adsmax 4.3 at pH 7.5

Equil at 1 h; Redlich-Peterson [73]

S 25 C: I, 11; 0.5; 0.9 cm2 area, 3.1 cmht,FL 1.5 mL/min

Pseudo-first order

C; C: – C: breakthrough column ads capacity4.16

W 25

Fly ash (class F, 9.1%CaO)

C; 10.1; I, 0–100; 450; F concentration in effluent reached0 mg/L after 120–168 h

– – [74]

S 20 40 cmht, 4.5 cmd, FL 2 mL/h

Fly ash (2.22% CaO) B; 2–9.5; E, 0–3; Ads max at pH 6.5; Langmuir ads max20 at pH 6.5

Pseudo-first order Langmuir [46]

S 30 20

Alum sludge (calcinedand uncalcined)

B; 3.5–8.8; E, 0–15; Ads max at pH 6; calcined higher adscapacity than uncalcined; Langmuirads max 5.39 at pH 6

Pseudo-first order;diffusion

– [48]

S 32 0.5–16

Spent catalyst B; 2–9.5; E, 0–40; Ads max (1.6 for initial F 20 mg/L,adsorbent 10 g/L) at pH 2, thencontinue to decrease. Max ads 28

Pseudo-first order;70% ads in 10 min

– [75]

S 25 10

ardous

ambmc

iBciTa

bsFSrtrSoDic

fSowwabecaioaN

(atanc

3

teflcprapa

ntu(r

P. Loganathan et al. / Journal of Haz

nion exchange resin which adsorbs anions by an ion exchangeechanism, whereas these anions had no effect on F adsorption

y a chelating resin which adsorbed F selectively by a H-bondingechanism. The preferential order of adsorption of anions by the

helating resin was reported to be F− > Cl− > NO3− > SO4

2−.Solangi et al. [34] studied the adsorption of F by a thio-urea

ncorporated amberlite resin in the presence of PO43−, Cl−, SO4

2−,r−, NO3

−, NO2−, HCO3

−, and CO32− at five times the molar con-

entration of F and observed that Br−, NO2−, and PO4

3− had littlenterference with F adsorption; the other ions had no interference.his was explained as due to the strong H-bonding of F with themide groups in the resin.

Multivalent metal oxides are known to adsorb F selectivelyy the ligand exchange specific adsorption mechanism. An alumludge containing a high percentage of Al, Ti and Fe oxides adsorbed

selectively in the presence of SO42− and NO3

− [48]. At 50 mg/L,O4

2− and NO3− reduced F adsorption from 85% to 40% and 62%,

espectively from a solution containing 20 mg F/L. In contrasto these anions, PO4

3− and selenate concentrations at 20 mg/Leduced F adsorption to 25% of F in solution. Based on these results,ujana et al. [48] proposed that the decreased order of competitionf anions for F adsorption to be PO4

3− ≥ selenate > SO42− > NO3

−.as et al. [55] also found that PO4

3− interfered more than SO42−

n F removal by a calcined Zn/Al layered double hydroxide (LDH)onsisting of Zn/Al oxide.

Kumar et al. [7] investigated the adsorption of F by a granularerric hydroxide in the presence of competing anions such as Cl−,O4

2−, BrO3−, NO3

−, CO32− and PO4

3−, each having concentrationsf 20– 100 mg/L with an initial F concentration of 20 mg/L. Thereas no significant influence of competing anions on F removalhen the adsorbent dose was 10 g/L, which was attributed to the

vailability of plenty of sorption sites. However, when the adsor-ent concentration was reduced to 5 g/L, PO4

3−, CO32−, and SO4

2−,ach at concentrations of 100 mg/L, this reduced the F adsorptionapacity by 35, 25, and 20% of F in solution, respectively. The othernions did not significantly reduce the F adsorption capacity. Sim-larly, Raichur and Basu [50] found that F adsorption by a mixturef naturally occurring rare earth oxides (oxides of La, Ce, Pr, Nd, Smnd Y) was not significantly affected by the presence of SO4

2− orO3

− in water at a concentration equal to that of F (100 mg/L).It can be concluded that the non-specifically adsorbing anions

e.g. nitrate, chloride) do not compete with F for adsorption ondsorbents that adsorb F using specific adsorption. Only anionshat are adsorbed by specific adsorption (e.g. phosphate, selenate,rsenate) compete with F for adsorption. When F is adsorbed byon-specific adsorption the non-specifically adsorbing anions canompete with F.

.3. Temperature

Temperature had no consistent effect on F adsorption. Adsorp-ion on many adsorbents increased with temperature showing anndothermic nature of adsorption (granular ferric hydroxide [7],y ash [46], calcined Mg/Al/CO3 LDH [54], LDH/chitosan [63], spentatalyst [75]). In contrast with many others it decreased with tem-erature showing an exothermic nature of adsorption (chelatingesin [30], alum sludge [48], calcined Zn/Al LDH [55], modifiedctivated carbon [59], LDH/chitin [65], geo-materials [68]). Tem-erature has also been reported to have no significant effect ondsorption by some adsorbents (trivalent cations/zeolite [40]).

The reasons for the differences in the effect of temperature wereot clearly stated in the studies. It may depend on the tempera-

ure range studied, the nature of the adsorbent, and the conditionssed in the studies. For example, at extremely low temperatures5, 10 ◦C) the rate of adsorption was reported to be low because theate of movement of F to the adsorption sites is low. Lai and Liu [75]

Materials 248– 249 (2013) 1– 19 11

reported that F adsorption on spent catalyst increased significantlywhen the temperature increased from 5 ◦C to 25 ◦C but very littlefurther increase in adsorption was observed when the tempera-ture rose to 50 ◦C. Sujana et al. [48] reported that the exothermicnature of adsorption of F on alum sludge existed because the risingtemperature increased the tendency for F to escape from the adsor-bent. Another reason given was an increase in thermal energy ofadsorbed F at higher temperatures, causing increased desorption.

3.4. Adsorption kinetics

Fluoride adsorption studies have shown that the rate of removalof F by adsorbents is high in the initial 5–120 min where generallymore than 90% of F is adsorbed, but thereafter the rate significantlylevels off and eventually approaches zero denoting the attainmentof equilibrium. This is because initially the adsorption sites arevacant and the F concentration gradient between solution andadsorbent surface is high. Subsequently the rate decreases becauseof the decrease in vacant sites. A fast rate of adsorption helps theadsorbent to treat large quantities of water [54]. A slow rate causesoperational, control, and maintenance problems in the adsorptionprocess of the filter bed [31].

The rate of F adsorption increases with an increase in concen-tration of adsorbent [54] and a decrease in initial F concentration[29,47,54]. It also depends on the structural properties of the adsor-bents and the interaction between F and the sites of adsorption.Adsorption kinetics has been described by pseudo-first order andpseudo-second order kinetic models and the diffusion model andrate constants have been determined (Table 1). These models havealso provided information on adsorption mechanisms.

4. Adsorbents

The effectiveness of F adsorption from drinking water by var-ious adsorbents, the methods used for the assessment, and thekinetics and equilibrium models that best explained the adsorp-tion process are presented in Table 1. Caution needs to be exercisedin comparing the adsorption capacities of adsorbents because ofthe inconsistencies in data presentation including differences inmethodology and parameters used in the studies (pH, temperature,F concentration range, competing ions, etc.). An ideal adsorbent thatcan be used to remove F must have the following characteristics:low-cost, a high F adsorption capacity, rapid adsorption of F, eas-ily regenerated after its removal capacity is exhausted and goodphysical characteristics (rapid water flow without filter clogging).

4.1. Metal oxides and hydroxides

Oxides and hydroxides, also called hydrous oxides or oxyhy-droxides, of trivalent and tetravalent metals such as Fe, Al, La,Mn, and Zr are used to remove both anionic and cationic con-taminants from water and wastewaters because of their strongability to adsorb these ions [37,77]. The predominant mechanismof adsorption of F on oxides and hydroxides is ligand exchange bythe formation of inner sphere complexes (specific adsorption) asdiscussed previously. Most metal oxides and hydroxides have theirPZC above the natural water pH of 7 (granulated ferric hydroxide pH7.5–8.0 [7], activated alumina pH 8.25 [36], �-alumina pH 8 [49]).Therefore, at the neutral pH of natural water, these adsorbents havepositive surface charge which is favourable for the adsorption of thenegatively charged F.

Of the oxides and hydroxides of metals, Al oxide, especially theactivated form (activated alumina) has been the most commonlyused adsorbent for the removal of F [9,29,36,37,42]. Activated alu-mina is produced by thermal degradation of aluminium hydroxide

1 ardous

tb([vato(1

ctt(niiFc

b[ato

rgbopTpoAFaeoro

pacwoabOLwMum1t

aArf2Nfi

2 P. Loganathan et al. / Journal of Haz

o obtain materials with high specific surface area and a distri-ution of micro- and macro-pores [78]. The specific surface aream2/g) of activated alumina has been reported to be 160 [49], 29729], and >200 [26]. The adsorption capacity of activated aluminaaries with the structure of the alumina. For example, �-Al2O3 has

much higher adsorption capacity than �-Al2O3 [42] and thereforehis form of activated alumina is commonly used for defluoridationf water. The maximum F adsorption capacity of activated �-Al2O3mg/g) has been reported to be 1.1 [36], 12.0 [41], 2.41 [47], and6.3 [49].

Aluminium oxides have poor adsorption capacity for F in acidiconditions because of their tendency to dissolve and form posi-ively charged AlF complexes (AlF2+, AlF2

+) which are repelled byhe positively charged surfaces of Al oxides at these conditionsPZC of Al oxides > pH 7) [49]. In alkaline conditions, Al oxides haveegatively charged surfaces which repel the negatively charged F

ons present at these pHs. Also, as stated previously, the increasen the concentrations of OH− in alkaline condition competes with− for adsorption. Therefore the optimum pH for F adsorption isonsidered to be near neutral pH [47,49].

Recent studies have shown that surface coatings of adsor-ents with other materials have enhanced the adsorption of F29]. Because manganese oxides have a high specific surface area,

micro-porous structure [79] and a high adsorption capacityowards anions [80], activated alumina was coated with thesexides and F adsorption capacities have been studied [29,36].

Maliyekkal et al. [36] compared the F adsorption capacity andate of adsorption of activated alumina (AA) and a granular man-anese oxide coated AA (MOCAA) at pH 7 and found that in theatch study, most of the adsorption was complete in 3 h in the casef MOCAA compared to 10 h for AA. The pseudo-first order andseudo-second order rate constants were also higher for MOCAA.he Langmuir adsorption capacity of MOCAA was 2.85 mg/g com-ared to 1.08 mg/g for AA. Maximum F adsorption was found toccur at a wider pH range of 4–7 for MOCAA compared to that forA (4–6). The column study also showed that MOCAA had higher

adsorption capacity. The superiority of MOCAA over AA in thedsorption of F was reported to be not due to the surface area differ-nce because the specific surface area of AA was 204 m2/g and thatf MOCAA was 170 m2/g. Maliyekkal et al. [36] suggested that theeason for this could be the increased zeta potential (surface charge)f the MOCAA, although no supporting data were presented.

Teng et al. [29] used a redox process to coat AA with an amor-hous MnO2. The granular MOCAA produced had a higher surfacerea (316 m2/g) compared to that for AA (297 m2/g) and a signifi-antly rougher surface with plenty of pores. Maximum F adsorptionas obtained at the pH range of 4–6. One of the two mechanisms

f F adsorption at a pH less than 6 was considered to be chemicaldsorption by a ligand exchange of surface OH groups in MOCAAy F in solution resulting in an increase of pH due to the release ofH−. The other mechanism was intra-particle diffusion of F. Theangmuir adsorption maximum at pH 5.2 was 7.09 mg/g whichas much higher than the value of 1.08 mg/g reported for AA byaliyekkal et al. [36] and values reported for many other gran-

lated adsorbents. For initial F concentrations less than 21 mg/L,ost of the adsorption was completed within 30 min compared to

0 h for AA. Faster adsorption by MOCAA was considered to be dueo the larger surface area and porous surface of this adsorbent.

Iron oxides are also known to have large capacity to removenions from water by mechanisms similar to those operating inl oxide adsorbents [81]. Kumar et al. [7] studied the adsorptiveemoval of F by a highly porous and poorly crystalline granulated

erric hydroxide (GFH) (�-FeOOH) with a specific surface area of50–300 m2/g, a PZC of pH 8, and a granular size of 0.32–2.0 mm.early 95% adsorption of F from solution was achieved within therst 10 min of agitation of 10–20 mg F/L with 10 g GFH/L at pH 6–7.

Materials 248– 249 (2013) 1– 19

The maximum F adsorption was observed at the pH range of 3–8.A sharp decrease of F adsorption was observed above the PZC pHof 8 as the GFH surface became more negatively charged causingelectrostatic repulsion of the negatively charged F− ions in additionto increasing concentration of OH− ions which competed with F−

for adsorption.Hydroxyapatite, the most abundant of phosphate minerals, has

been used for defluoridation of water. Fan et al. [52] compared theF adsorption capacities of several natural minerals and found thatthe adsorption capacities at pH 6 followed the order: hydroxyapa-tite > fluorspar > quartz activated using ferric ions > calcite > quartz.The highest F adsorption capacity of hydroxyapatite was explainedas owing to F exchanging with a OH group at the surface and insidethe apatite mineral. The F adsorption on the other minerals wasdeemed to be a surface adsorption process.

It is evident from the literature review that Fe and Al oxides andhydroxides are the commonly used adsorbents for defluoridation.They have a moderate level of F adsorption capacity (1–16 mg/g)and if locally available at low cost, can potentially be employed inrural areas, especially in developing countries.

4.2. Layered double hydroxides

The majority of clay minerals such as kaolinite, mica, mont-morillonite, vermiculite and zeolite, carry predominantly negativecharges and therefore adsorb very small amounts of anions. Lay-ered double hydroxide (LDH) or hydrotalcite (HTlc) is another typeof clay mineral, but has positive charges and therefore adsorb sig-nificant quantities of anions and oxyanions (e.g. fluoride, arsenite,arsenate, chromate, phosphate, selenite, selenate, nitrate, etc.) andmonoatomic anions (e.g. fluoride, chloride) from aqueous solu-tions [82]. Structurally, LDHs are composed of positively chargedbrucite-like sheets compensated by a large number of exchange-able charge-balancing anions in the hydrated interlayer regions[55,56,82]. Charges can also be produced by the ionisation of theOH groups at the surface of the LDH particles. The PZC of LDHs is inthe region of pH 9–12 [82] and therefore at the neutral pH of naturalwater, LDHs carry positive charges and act as anion exchangers.

Calcined LDHs have higher F adsorption capacities than theuncalcined LDHs [53]. The optimum temperature of calcinationis generally considered to be 450–500 ◦C [53–56]. Lv et al. [54]reported that the F adsorption capacity of a Mg/Al LDH increasedfrom 65 to 70 and then to 80 mg/g when the calcination tem-perature was increased from 200 ◦C to 400 ◦C and then to 500 ◦C,respectively, but decreased to 62 and 50 mg/g when the tempera-ture was increased to 600 ◦C and 800 ◦C, respectively. Wang et al.[53] suggested that the increase in F adsorption capacity of LDH asa result of calcination was due to the higher specific surface area,porosity, and surface reactivity of the Mg/Al oxide produced by cal-cination. Another reason reported was the incorporation of F intothe structure of Mg/Al oxide resulting in the formation of the orig-inal structure of the LDH. The decrease in F adsorption capacityat calcining temperatures higher than 500 ◦C was considered to bedue to the transformation of the Mg/Al oxides into a spinel structurethat does not exhibit the property of LDH structural reconstruction[54].

The adsorptive property of LDH depends on the metallic con-stitution of the LDH structure. Lv et al. [54] showed that the Fadsorption capacity of calcined Mg/Al LDH was higher than thatof calcined Ni/Al LDH and calcined Zn/Al LDH because of the higheratomic weights of Ni and Zn compared to Mg. Among the calcinedMg/Al LDHs, the one having Mg/Al molar ratio of 2 was found to

have the highest F adsorption capacity. The maximum F adsorptioncapacity of 213 mg/g was obtained at pH 6. This adsorption capacityis the highest of all adsorbents listed in Table 1. However, Das et al.[55] reported a lower Langmuir adsorption capacity of 17 mg/g at

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P. Loganathan et al. / Journal of Haz

H 6 for a calcined Zn/Al LDH. The F adsorption data of Wang et al.53] on a calcined Mg/Al LDH did not fit the Langmuir adsorption

odel but the amount of F adsorbed continued to increase withquilibrium F concentration with an adsorption capacity of 35 mg/gt the highest tested F equilibrium concentration of 70 mg/L. Theate of F adsorption on calcined LDH is variable [53,55,56].

The literature review reveals that LDHs can have very high adsorption capacity (17–213 mg/g) if calcined to 500 ◦C. Thedsorption capacity varies depending on the type and proportion ofhe metals in the LDH structure. Because LDHs have high adsorptionapacity, they are useful adsorbents for defluoridation of watersith high F concentrations.

.3. Ion exchange resins and fibres

Ion exchange resins and fibres are an important class of adsor-ents used to remove anionic and cationic pollutants from waternd wastewater. Their framework or matrix consists of irregular,acromolecular, three dimensional network of hydrocarbon chain

32]. The cation exchange resins and fibres have negatively chargedunctional groups whereas the anion exchange resins and fibresave positively charged functional groups such as −NH3

+, NH2+,

N+, S+. Therefore the cation exchangers adsorb cations and thenion exchangers adsorb anions such as F−. The cation exchangersan also be made to adsorb anions if they are impregnated with pos-tively charged metallic cations that have strong affinity for anions39,57].

Ku et al. [57] used an Al incorporated cation exchange resinAmberlite IR-120) to remove F from water and found that the Lang-

uir maximum F adsorption capacity at pH 4 was 4.6 mg/g. Withinhe pH range of 4–9 tested, F adsorption emerged as the greatestt pH 5–7. Metals other than Al have also been incorporated intoation exchangers to enhance the F adsorption capacity. For exam-le, Luo and Inoue [39] compared the F adsorption capacities of theation exchange resin, Amberlite 200 CT modified by the incorpo-ation of a number of trivalent metal ions, La, Ce, Y, Fe, and Al. The

adsorption capacity of the metals incorporated resins in the pHange of 4–7 was in the order: La ≤ Ce > Y > Fe ∼ Al. More than 80%f F was adsorbed by 20 g La/L resin from a 15 mL solution contain-ng 15 mg F/L. The resin without any metal incorporation adsorbednly about 5% of F from the solution.

The majority of anion exchange materials are not effective indsorbing F from natural water containing other anions because− affinity to anion exchange materials is the least of all anionscitrate > SO4

2− > oxalate > I− > NO3− > CrO4

2− > Br− > SCN− > Cl− >ormate > acetate > F− [32]). Consistent with this order of affinity,

strong base anion exchanger was shown to effectively adsorbO3

−, Br−, and SO42−, but not Cl− and F- [83].

The F adsorption capacity of anion exchangers can be enhancedy modifying the adsorbing sites on the anion exchangers. Gener-lly, anion exchange resins impregnated with chelating agents thatan form H-bonding with F are used. Meenakshi and Viswanathan30] compared the F adsorption capacity of a chelating resin hav-ng sulphonic acid functional group and an anion exchange resinnd showed that 1 g of the chelating resin adsorbed 95% of F from

50 mL solution containing 3 mg F/L in 40 min compared to 65%dsorption by the anion exchange resin.

Solangi et al. [34] reported that 100 mg of the ion exchange resin,mberlite XAD-4 modified by incorporation with thio-urea bindingites removed 90% of the F from a 10 mL solution containing 16 mg/L compared to 30% by the unmodified resin at pH 7. The higher

adsorption capacity of the modified resin was explained as being

ue to H-bonding between the amide groups in thio-urea and F. Theangmuir maximum F adsorption capacity of the modified resin atH 7 was much higher than the F adsorptive capacities of manyther adsorbents in the literature (Table 1). Interference of other

Materials 248– 249 (2013) 1– 19 13

co-existing anions present at the concentration ratio of 1:5 (F:otherions) was insignificant. Therefore the authors concluded that themodified resin can be effectively used for defluoridation of water.

Fibrous adsorbents, due to their physico-chemical structure,generally have rapid rates of adsorption of ions. If their adsorp-tive capacity and affinity can also be enhanced, they can be a usefulclass of adsorbents for removal of ions from water. Ruixia et al. [31]introduced functional groups into a polyacrylonitrile ion exchangefibre and studied its F adsorption behaviour. They found that the Fadsorption reached equilibrium in 5 min with 40% of F adsorptionfrom a 500 mL solution containing 5, 34, and 50 mg/L of F, As(V), andP, respectively when 1 g fibre modified with functional groups wasadded. The adsorption by the unmodified fibre was less than 5% ofsolution F. The rate of adsorption was considered to be faster thanthat observed in many other adsorbents. Maximum adsorption wasobtained at pH 3. The mechanism of adsorption was considered tobe ion exchange.

It could be concluded from this review that ion exchange resinsand fibres have poor F adsorption capacities but the adsorptioncapacities and selectivity for F adsorption in the presence of otherions can be significantly increased (up to 61 mg/g) by surface modi-fication of the adsorbents by loading with organic functional groupsand metals. These adsorbents are relatively expensive and thereforethey are useful only for water treatment in industrial countries.

4.4. Zeolites

Because zeolites have negative surface charges at all pH val-ues, they have high adsorption capacity for cations, but have lowadsorption capacity for anions. Nonetheless their adsorption capac-ity for anions can be increased by modifying the zeolite surfacewith cationic surfactants or multivalent metallic cations [85]. ForF adsorption, only metallic cations incorporated with zeolites havebeen used. There appear to be no studies conducted on usingsurfactant- or organic compounds-modified zeolites on defluori-dation.

Onyango et al. [40] modified the surface of a synthetic zeoliteby exchanging Na+ in the zeolite with Al3+ and La3+ to create activesites for F adsorption. The introduction of Al and La opened upthe pores in zeolite leading to increased porosity. The Langmuiradsorption capacity is slightly larger for La-zeolite than for Al-zeolite (Table 1). However, within the F concentration range studied(10–80 mg/L), Al-zeolite had twice as much adsorption capacity asthe La-zeolite. The adsorption capacities obtained for these zeo-lites were reported to be much higher than many other adsorbentsincluding the commonly used activated alumina (Table 1). This sug-gested that the mechanism of adsorption of F onto Al-zeolite wasmostly by a chemical adsorption process (ligand exchange) andadsorption onto La-zeolite was mostly by a physical adsorptionprocess (coulombic attraction).

The Na-zeolite had no PZC because it was negatively chargedat all pHs between 3 and 13 and therefore the adsorption of F− ispoor due to charge repulsion. In contrast, the Al-zeolite and theLa-zeolite had PZC at pHs of 8.15 and 4–5.25, respectively indicat-ing that below these pHs these zeolites were positively charged. Inaccordance with the surface charge on La-zeolite the F adsorptionwas the highest at pH 5; and decreased above and below this pH. Inthe case of the Al-zeolite, the increase of pH increased F adsorptionwith a plateau forming above pH 5. The F adsorption on Al-zeoliteat pHs above 5 did not decrease as observed for La-zeolite up tothe highest pH 9 tested because F was reported to be adsorbed bychemical adsorption. The authors stated that, over all, Al-zeolite

was superior to La-zeolite for defluoridation.

In a subsequent study by Onyango et al. [41] on F adsorptionby four types of Al pre-treated low-silica synthetic zeolite it wasobserved that F adsorption on all four types increased from pH 2 to

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4 P. Loganathan et al. / Journal of Haz

and remained constant up to pH 8, before decreasing from pH 8o 11. The pH effect on F adsorption was explained by the followingeactions.

HF � H+ + F− (low pH)Zeolite − AlOH+

2 + F− � Zeolite − AlF + H2O (neutral pH)Zeolite − AlOH + F− � Zeolite − AlF + OH− (neutral pH)Zeolite − AlOH + OH− � Zeolite − AlO− + H2O (high pH)

At low pH, part of the F is removed as weakly ionised HF andherefore F adsorption is reduced (first equation). Another reasonould be that some of the F is complexed to Al solubilised at lowHs to form AlF + complex which has low tendency to adsorb onhe positively charged adsorbent at the low pHs. At high pH, theegatively charged adsorbent (last equation) repels the negativelyharged F− as well as increased competition of OH− for adsorption.

Samatya et al. [23] using a natural zeolite from Turkey pre-reated with La, Al, and Zr to study the removal of F from tap waterpiked with NaF. They showed that the F adsorption capacities athe equilibrium F concentration range of 0–12 mg/L were largestor Zr-zeolite and smallest for La-zeolite. The values for the La andl zeolites were lower than the values reported by Onyango et al.

40], probably because the metal loadings were much lower andhe zeolite used was natural compared to the synthetic zeolite ofnyango et al. [40]. However, all three metal zeolites removed 95%f F from an aqueous solution containing 2.5 mg F/L at an adsorbentose of 6 g/L.

In contrast to the above studies which showed high F removalapacities of multivalent metal incorporated zeolites, Diaz-Navat al. [85] found that La and Eu treated natural zeolite from Mexicoad only slightly higher F removal capacities than when this zeoliteas treated with Na and Ca. The weight percentages of the metals

n the zeolite were 1.7, 3.20, 0.32, and 1.49 for Na, Ca, La, and Eu,espectively.

The review revealed that zeolites on their own had very lowdsorption capacities for F but when they were loaded withultivalent metallic cations they produced moderate to very

igh F adsorptive capacities (2–45 mg/g). The adsorptive capacityepends on the type and amount of metal loading. Zeolites can alsorovide good physical properties for practical use.

.5. Carbon materials

Activated carbon (AC) is an important carbon material com-only used as an adsorbent for the removal of a wide range

f aquatic pollutants due to its exceptionally high surface area500–1500 m2/g), highly developed internal microporosity, pres-nce of a range of functional groups, low cost and ready availability86,87]. However, it displays poor adsorption capacity towardsnionic pollutants because of its low PZC (pH 1.6–3.5 [88]).

The amount of anions adsorbed onto AC depends on the poreize distribution because the adsorption occurs mainly in the pores.be et al. [60] reported that F adsorption capacity increased with

he specific surface area in 11 out of the 12 carbonaceous materialstudied by them. The adsorption of F on bone char did not fit thisattern because the mechanism of adsorption was different fromhe rest. In bone char, F is adsorbed chemically by ligand exchangeith the OH group in the hydroxyapatite compound [1]. The inef-

ectiveness of carbon in adsorbing F was also shown by Srimuralit al. [70] who found that lignite coal and char fines, a by-productbtained during the making of coke, adsorbed only 7.9 and 19% of F,espectively from a 50 mL solution containing 5.0 mg F/L and when00 mg of adsorbents were added and mixed for 5 h. In comparison,

entanite clay removed a much larger percentage of 33% of solution.

A new type of carbon material called aligned carbon nanotubesACNT) made up of needle-like cylindrical tubules of concentric

Materials 248– 249 (2013) 1– 19

graphitic carbon capped by fullerene-like hemispheres was devel-oped as a promising adsorbent material for the adsorption ofcontaminants in water [12]. Li et al. [12] found that ACNT had higherF adsorption capacity than AC (Table 1), despite the ACNT havinga much lower specific area and pore volume than the AC. Li et al.[12] stated that though the adsorption capacity of the ACNT washigh, the cost too was high which may limit its full commercialutilisation.

The surface of the carbon particles can be modified to improvethe F adsorptive properties. Such modifications have been broughtabout by creating new functional groups which have strong affin-ity towards F. Daifullah et al. [59] modified the structure of an ACproduct by steam pyrolysis of rice straw and oxidised the productusing HNO3, H2O2, and KMnO4. Of these treatments, the materialobtained by KMnO4 oxidation produced the highest F adsorption.The adsorption mechanism was considered to be ligand exchangeof OH groups on the carbon surface with F. The presence of MnO2 onthe carbon surface caused by reduction of the KMnO4 may have alsoparticipated in the removal of F. Thermodynamic studies showedthat the adsorption was chemical in nature. The Langmuir adsorp-tion maximum was 15.9 mg F/g at the natural pH of water. Thisvalue was considered to be higher than the values reported in theliterature for many other adsorbents.

Ramos et al. [42] studied the F adsorption behaviour of Al-impregnated AC (AlAC) produced from coconut shells followed bycalcination to 300 ◦C. The AlAC had a F adsorption capacity morethan four times higher than the plain carbon at an equilibriumconcentration of 2–8 mg F/L. The Langmuir adsorption capacity ofAlAC (1.07 mg/g) at natural water pH, though higher than the valueobtained for plain carbon (0.49 mg F/g), was lower than the valuesobtained for some other adsorbents especially activated alumina.The F adsorption reached its maximum at pH 3 because the PZCof the AlAC was 4.1–4.8 (above this pH, AlAC had negative surfacecharges which repel F−). However, there was appreciable adsorp-tion of F at pH 6–7 (0.6 mg F/g at the equilibrium F concentrationof 6 mg F/L) due to chemical adsorption of F onto Al which does notinvolve electrostatic attractive forces.

In a later study, Li et al. [61] reported that the adsorption capacityof aligned carbon nanotubes (ACNT) at pH 6 increased from 2.3 mgF/g to 9.6 mg F/g when Al was loaded onto ACNT and calcined to450 ◦C at an optimum Al2O3 rate of 30% by weight. The adsorptioncapacity fell to less than 6 mg F/g when the loading rate was 20 and40%. The maximum adsorption capacity was obtained at pH 6.0–9.0.The mechanism of F adsorption at pH less than the PZC of 7.5 wasconsidered to be by electrostatic attraction of negatively chargedF− ions onto the positively charged adsorbent as well as by ligandexchange of the OH group by F. At the PZC and slightly above thispH the adsorption was reported to be mainly by ligand exchange.At pHs much higher than the PZC (pH > 9), OH− ions in solutioncompeted with F− ions for adsorption and therefore F adsorptiondecreased.

It can be concluded from the review that AC is a poor adsorbentfor F. However, its adsorption capacity can be increased to moder-ate levels (up to 16 mg/g) by chemical modification of the carbonsurface. Specialised carbon materials such as ACNT have higheradsorption capacity than AC. Their adsorption capacity can also beincreased by surface modification but this will incur additional cost.

4.6. Natural materials

Several natural inorganic materials (soils, clays, minerals, andbuilding materials) have been used in the defluoridation of water.

The adsorption capacities of these materials have been discussedin previous reviews (17, 26, 27) and also presented in Table 1.

Bioadsorbents contain a variety of functional groups such as car-boxyl, imidazole, sulphydryl, amino, thioether, phenol, carbonyl,

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P. Loganathan et al. / Journal of Haz

mide, and hydroxyl moieties capable of adsorbing the pollutants,specially the metal cations [78]. Considering this feature of bioad-orbents, metals having strong affinity towards F were loaded ontoelected bioadsorbents (Zr4+ on collagen fibres [13], La3+ on geletin44]) for enhancing the adsorption of F from drinking water.

Chitin and chitosan have been shown to constitute anotherffective group of bioadsorbents for removing a range of aquaticollutants due to their low cost and high contents of amino andydroxyl functional groups [89]. Chitin is a polysaccharide found

n a wide range of organisms but most commonly extracted fromhellfish processing waste, while chitosan is a copolymer of glu-osamines derived from chitin by deacetylation in hot alkali [67].ahli et al. [22] reported that the adsorption of F on chitosan from

synthetic water was rapid and reached a maximum at 5 min.ncrease of pH from 2 to 6 increased adsorption followed by aecrease up to pH 10. The F adsorption from a brackish waterontaining 3.25 mg F/L and several other anions at much higheroncentrations (Cl− 1083 mg/L, SO4

2− 215 mg/L, HCO3− 171 mg/L)

howed that the selectivity of chitosan towards various anions wass follows: F− > HCO3

− > NO3− > Cl− > SO4

2−. In a batch adsorptiontudy of 15 min duration the F concentration fell to 1.82 mg/L, butith a successive batch adsorption the concentration reduced to

.9 mg/L which was lower than the WHO standard of 1.5 mg F/L.Ma et al. [66] prepared a magnetic-chitosan adsorbent by co-

recipitation of chitosan with Fe2+ and Fe3+ salts and used it toemove F from a synthetic F solution in a batch study. The adsor-ent was shown to have a higher F adsorption capacity (Langmuirdsorption capacity 22.49 mg F/g) than a commercial activated alu-ina (Table 1). Adsorption capacity was highest in the pH range

–9. The FTIR data of the adsorbent showed that the main groupsnvolved in the adsorption process were amine and iron oxidesn the magnetic-chitosan. The main advantage of having magneticroperties in the adsorbent is to easily separate the adsorbent fromhe treated water after its use using an external magnetic field sohat the adsorbent can be reused.

A research group in India has reported that the F adsorp-ion capacity of nano-hydroxyapatite (nHAp) can be increased byncorporating chitin or chitosan onto nHAp [63–65,90]. In a batchxperiment conducted by Sundaram et al. [65], it was observedhat mixing 0.25 g each of chitosan, nHAp, and nHAp/chitosan com-osite with 50 mL solution of NaF containing 10 mg F/L, removed.05, 1.30, and 1.56 mg F/g adsorbent, respectively after 30 min ofhaking the suspensions. Defluoridation of water from a fluoride-ndemic village using nHAp-chitosan composite was found to beigher compared to using nHAp. Similar results were obtained from

subsequent study where a composite of nHAp and chitin was used64]. However, the adsorption capacity of this composite was higherhan that of the composite made with chitosan (Table 1).

The review revealed that although the majority of naturalaterials have low effectiveness in removing F (<1 mg/g), they

re inexpensive and consequently can be used in less developedountries, especially in rural areas. Furthermore, their F removalapacities can be increased through the process of surface modifica-ion, such as loading with multivalent metallic cations, as reportedreviously for other adsorbents.

.7. Industrial by-products

Several types of industrial by-products have been used for thedsorptive removal of pollutants including fluoride from water27,28,89,91,92]. Most of the adsorbents fall into the categories ofy-products from the mining industry (e.g. red mud), steel industry

e.g. slag materials), and power plant industry (e.g. fly ash).

Cengeloglu et al. [4] investigated the adsorption of F from aque-us solutions using a red mud containing 18.7% Al2O3, 39.7% Fe2O3,nd 14.5% SiO2 with (ARM) and without (RM) activation by 20%

Materials 248– 249 (2013) 1– 19 15

HCl treatment in a batch study. The amount of F removed from asolution (pH 5.5) containing 3.8–28.5 mg F/L by ARM at a dose of0.2 g/50 mL was at least three times higher than that by RM at thesame dose. The main mechanism of F removal was considered tobe that of specific adsorption caused by ligand exchange of F andOH groups on the metal oxide surfaces in red mud.

Red mud particles are too fine for use in filter beds becauseof their poor hydraulic conductivity. This problem can be over-come either by mixing red mud with coarse-size particles such assand [93] or by granulation [94] before use. Tor et al. [72] used themethod of Zhu et al. [94] for the granulation of red mud by mixing itwith fly ash, sodium carbonate, quicklime and sodium silicate. Themaximum F adsorption (2.2 mg/g) from a NaF solution containing15 mg F/L and granulated red mud (GRM) of 5 g/L was obtained ata pH of 4.7 in the pH test range of 2.5–7.3. The adsorption capacitydetermined from a column study was 2.05 mg F/g for a flow rate of2 mL/min compared to 0.644 mg F/g obtained in the batch study atthe same GRM dosage (5 g/L) and initial F concentration (5 mg/L).

Fly ash is a major by-product that is produced from the com-bustion of coal in power stations. It consists of fine and powderymaterials (1.0–100 �m) made up of a mixture of amorphous andcrystalline alumina-silicates and several compounds of Si, Al, Fe, Ca,and Mg [95] and therefore a good candidate material for F adsorp-tion from water. Chaturvedi et al. [46] reported that the adsorptionof F from water by a fly ash containing 56.0% SiO2, 25.9% Al2O3,2.22% CaO, and 1.26% Fe2O3 had a maximum Langmuir adsorptioncapacity of 20 mg/g at pH 6.5. The F adsorption increased from 79to 94% when the pH of the F solution (10 mg/L) increased from2.0 to 6.5 and then decreased with further increase in pH up to9.5. However, Nemade et al. [62] reported that F adsorption by afly ash decreased continuously from pH 2 to 12. The difference inthe pH effects between the two studies could be due to the differ-ence in chemical characteristics of the fly ashes and experimentalconditions used.

In a column study, Piekos and Paslawska [74] found that an alka-line fly ash (pH ≥ 10, 9.1% CaO) effectively removed F from watercontaining 1–100 mg F/L when the solution was passed through acolumn (400 mm length) packed with 450 g fly ash at a flow rateof ≤2 mL/h. Complete adsorption of F was obtained after 120 h. TheF adsorption mechanism was suggested to be chemical binding ofF onto Ca(OH)2 and physical adsorption onto the residual carbonparticles in the fly ash.

Steel industry by-products such as blast furnace slag, electricarc furnace slag, basic oxygen furnace slag and converter slag havebeen shown to adsorb many pollutants, especially phosphate, fromwater due to the presence of high contents of Ca, Fe, Al, Mg, and Mnoxides [28]. However, limited studies have been conducted in theiruse in defluoridation of water [16,27]. Islam and Patel [71] showedthat a basic oxygen furnace slag (BOFS) containing 46.5% CaO, 16.7%Fe oxides, and 13.8% SiO2 had a good capacity to adsorb F. Ther-mal activation of the BOFS by heating to 1000 ◦C for 24 h increasedthe porosity and surface area leading to increased F adsorption.The percentage adsorption of F (initial F concentration 10 mg/L,0.5 g/100 mL adsorption dosage) was higher for the thermally acti-vated BOFS (TABOFS) (93%) than for the BOFS (70%). Unlike thecase with most adsorbents, increase of pH from 2 to 10 increasedF adsorption on TABOFS. The increased adsorption at high pH val-ues was considered to be due to F adsorption onto Ca(OH)2 in theslag. Presence of other anions reduced F adsorption in the order,PO4

3− > HCO3− > CO3

2− > SO42− > Cl− > NO3

−.Large quantities of alum sludge, a waste product, are generated

during the manufacture of alum from bauxite [48] and kaolin [96]

by the sulphuric acid process in many countries. As this sludgemainly consists of oxides of Al, it has a high adsorption capac-ity to remove many aquatic pollutants. Sujana et al. [48] useda sludge (47.2% Al2O3 and 20.7% TiO2) produced from an alum

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6 P. Loganathan et al. / Journal of Haz

anufacturing plant using bauxite raw material for the removalf F from water. The sludge after washing with water and calciningo 400 ◦C for 3 h adsorbed more than twice the amount of F thatas adsorbed by the uncalcined sludge from a solution contain-

ng 10 mg F/L of pH 6 (adsorbent dosage 20–300 mg/g). The higherdsorption capacity of the calcined sludge was considered to be duehe increased surface area produced by calcination. The F was con-idered to be specifically adsorbed according to a two-step ligandxchange mechanism. Within a pH range of 3–9, the maximum Fdsorption was observed at pH 6; and at both above and belowH 6 the adsorption decreased, as also reported for activated alu-ina by others. The presence of other anions in solution reduced F

dsorption according to the order, PO43− > silicate > SO4

2− > NO3−.

In contrast to F adsorption on sludge produced during the man-facture of alum from bauxite discussed in the previous paragraph,igussie et al. [96] reported that F adsorption by a sludge produceduring the manufacture of alum from kaolin raw material contain-

ng primarily SiO2 did not increase when the sludge was heatedo 300 ◦C. At temperatures above 300 ◦C up to 700 ◦C, the adsorp-ion decreased. Also, F adsorption was nearly constant from pH 3o 8 and significantly decreased above pH 10. The difference in thedsorption behaviour of the two types of sludge may be due to theifferences in the physico-chemical characteristics of the sludgeypes. However, the F adsorption capacities of the sludge in the twotudies were not too different. The presence of HCO3

− at concen-rations higher than those of F, decreased F adsorption efficiency,hile other anions (PO4

3−, SO42−, Cl−, NO3

−) had no significantffect [96].

Lai and Liu [75] used a by-product of the petrochemical industryalled spent catalyst consisting mainly of porous silica and aluminaspecific surface area 130 m2/g, PZC 5.2) to remove F from aque-us solutions. In a batch adsorption experiment on pH effect, theybserved that F adsorption decreased with increase in pH from 2 to

.5, with a plateau in the pH range of 4–7. The maximum adsorp-ion capacity at pH 4 was found to be 28 mg F/g at 25 and 50 ◦C.his value was considered to be comparable to that of activated alu-ina. Because the activation energy calculated using the Arrhenius

able 2hermodynamic parameters for fluoride adsorption by different adsorbents.

Adsorbent �G◦ (kJ/mol)Temperature (◦C)

Basic oxygen furnace slag −0.38 −0.72

25 35

LDH/chitosan composite −6.81 −6.74

30 40

Granular ferric hydroxide −39.2

25

Geomaterials −5.8 to −15.9

32 to 62

Nano-hydroxyapatite/chitin composite −7.1 −7.0

30 40

Nano-hydroxyapatite/chitosan composite −3.8 −3.4

30 40

Modified activated carbon −3.6 −0.9

25 45

Waste carbon slurry −25.4 −26.5

25 35

Chelating ion exchange resin −5.7 −5.5

30 40

Calcined Zn/Al LDH −25.0 −25.6

30 40

Fly ash −1.3 −1.6

30 40

Materials 248– 249 (2013) 1– 19

equation was very low (3.2 J/mol) the mechanism of adsorptionwas considered to be non-specific adsorption involving coulombicforces.

Many of the industrial by-products are wastes that need to bedisposed of. Their beneficial use can save disposal cost, preventenvironmental pollution arising from the disposal sites, and releasedisposal lands for alternative uses. In this respect, though thesematerials have low F adsorption capacities, they are cost-effectiveand therefore can potentially be used in developing countrieswhere cost of operation is a major factor in the choice of adsor-bents. Furthermore, many of these materials have been shown tobe good adsorbents for other aquatic pollutants and therefore itis possible to simultaneously remove many pollutants using thesematerials.

5. Adsorption thermodynamics

The strength and spontaneous nature of adsorption and infor-mation on whether the adsorption process is exothermic orendothermic are provided by determining thermodynamic param-eters such as changes in Gibbs free energy (�G◦), enthalpy(�H◦), and entropy (�S◦). These thermodynamic parametersare calculated from data on adsorption at different tempera-tures using standard thermodynamics equations (see references inTable 2).

A negative value for �G◦ indicates spontaneous and ther-modynamically favourable adsorption, while a negative �H◦

value indicates an exothermic adsorption process. An exother-mic reaction means that the amount of adsorption decreases withincreasing temperature. An endothermic reaction (positive �H◦

value) has been explained as due to enlargement of pore sizesand/or activation of the adsorbent surface [97,98]. Very low �H◦

values are generally associated with physical adsorption and very

high values with chemical adsorption [99]. However, no definitevalue for distinguishing the two forms of adsorption exists. �G◦

has also been used to distinguish between the two forms of adsorp-tion. For example, Meenakshi and Viswanathan [30] reported

�H◦ (kJ/mol) �S◦ (J/mol/◦C) Reference

−1.94 22.7 8.4 [71]45

−6.72 8.23 0.1 [63]50

7.34 15.8 [7]

−0.1 to −30 −3 to −10 [68]

−6.850 8.6 1.6 [65]

−3.450

0.9 −48.6 −15.1 [59]55

−27.6 7.3 11 [73]45

−5.6 −2.4 4.2 [30]50

−26.3 −5.7 6.4 [55]50

−2.0 6.5–6.9 17.1 [46]50

P. Loganathan et al. / Journal of Hazardous Materials 248– 249 (2013) 1– 19 17

Table 3Fluoride desorption and adsorbent regeneration.

Adsorbent (batch (B), column (C) method used) Desorption/regeneration reagent (BV,bed volume; d and ht, column internaldiameter and height; FL, flow rate

Results Reference

Amberlite modified with thio-urea (ATU) (B) and (C) (B) 100 g F-loaded ATU shaken 40 minwith 5 mL HNO3, H2SO4, HCl, NaOH(0.01, 0.001 M), 25 ◦C. (C) 2 g ATU,0.8 mmd, 50 mmht, FL 1 mL/min, 25 ◦C

(B) maximum desorption (99.9%) with0.01 M HCl. Next best 0.01 M H2SO4

(77.6%). (C) maximum recovery 92.3%(0.01 M HCl), 35 mL volume used

[34]

Manganese oxide coated alumina (MOCA) (B) 0.25 g MOCA adsorbed with F, 50 mLwater with pH adjusted to 3.0–13.5,shaken 24 h, 25 ◦C

85% desorption at pH 12. Little Fdesorbed at pH < 10

[29]

Granular red mud (GRM) (C) 10 g GRM, 0.635 cm2 cross sectionalarea, 10 cm ht, 10 mL 0.2 M NaOH, FL1 ml/min, 25 ◦C

Adsorption (mg/g) decreased from 2.05to 0.82 in 4 regeneration/adsorptioncycles. %desorption from 87 to 46

[72]

MOCA and activated alumina (AA) (C) 0.028 md, 0.1, 0.2 and 0.3 m ht, 2.5%NaOH. 3 cycles ofregeneration/adsorption

94.6%F desorbed for MOCA, 91%F forAA in Ist cycle. 2nd cycle no reductionin adsorption, 3rd cycle 5% reduction

[36]

Al treated zeolite (B) 0.1 g F loaded adsorbent, 50 mL aceticacid, water, NaHCO3, NaOH shaken for1d. 4 desorption steps. 25 ◦C

NaHCO3 most effective(63%desorption), Acetic acid leasteffective (5%). 4 zeolites with NaHCO3desorption, 1st step 61–67%, total75–95%

[41]

AA (C) 50.8 mm d, 550 mm ht, repeated cyclesof regeneration/reactivation

Loss of 2% F uptake capacity after 5cycles of regeneration

[47]

Zr loaded collagen fibre (B) 0.1 g, 100 mL water with pH adjustedwith NaOH, HNO3, 0.5 h shaking, 30 ◦C

pH < 9 very little desorption, pH 11.597% desorption

[13]

La loaded gelatin (B) 1st wash 1 M NaOH, then water wash oracid wash with 1:1 HNO3 to neutral pH

After adsorption/regeneration processfor 3 times, maximum adsorptioncapacity reduced from 98.5 to 82.3%

[44]

Calcined Zn/Al LDH (B) NaOH (0.001–0.04 M), 1 g adsorbenttreated with maximum F/L of NaOH for6 h, 30 ◦C.

F desorption increased from 5.12 to8.55 mg/g (100%). Desorptionincreased with NaOH concentration

[55]

Ion exchange fibre (B), (C) (B) 0.1 g adsorbent F loaded, mixedwith 25 mL 0.1 M NaCl, NaOH, HCl for2 h, 25 ◦C. (C) 1.5 cmd, 5 cmht, elutedwith 0.5 M NaOH, FL 1 mL/min, 25 ◦C

(B) Desorption 20% for HCl, 60% forNaCl, 80% for NaOH. (C) 100% desorbedwith 5 mL NaOH

[31]

Rare earth oxides (B) 0.2 g adsorbent, 100 mL water, pHadjusted, shaken for 30 min. 29 ◦C

pH < 6, F desorbed to ∼0. pH ∼ 12 > 95%.Regeneration decreased adsorption98–91%

[50]

Activated alum sludge (B) F loaded sludge 4 g/L, water pHadjusted, 4 h shaking, 30 ◦C

pH 2–7 desorption ∼ 0. >pH 7desorption increased, highest at pH 12

[48]

La loaded silica gel (C) Adsorbent loaded with F eluted with). 1 cm

Column was regenerated [43]

taassa

pontooiaTtt

6

aee

dilute NaOH (pH 8.50.5 mL/min, 20 ◦C.

hat �G◦ values of up to −20 kJ/mol were indicative of physicaldsorption, while values less than −40 kJ/mol involved chemicaldsorption. Positive �S◦ values indicate good affinity of adsorbedpecies towards the adsorbent and increased randomness at theolid-solution interface associated with structural changes at thedsorption sites during the adsorption process [99,100].

All studies, except the one on F adsorption at the highest tem-erature of 55 ◦C on activated carbon modified by H2O2/KMnO4xidation [59] presented in Table 2 had negative �G◦ values. Theegative values indicate good affinity of F for the adsorbents andhe adsorption process was spontaneous. The small positive valuebtained for the activated carbon was attributed to the disruptionf the MnO2 formed at the carbon surfaces and increase F solubil-ty at high temperature [59]. The �S◦ values were mostly positive,gain indicating a strong affinity of F towards the adsorbents.he �H◦ values, however, were positive and negative indicatinghe endothermic and exothermic nature of adsorption, respec-ively.

. Fluoride desorption and adsorbent regeneration

A suitable adsorbent for F removal should not only have high Fdsorption capacity and cost-effectiveness but also be amenable toasy desorption of the adsorbed F and capable of efficient regen-ration for multiple reuse of the adsorbent. Also, the desorbing

d, 9 cm h, FL

agent should not produce any damage to the adsorbent that causesa reduction in its adsorption capacity. Only adsorbents that canbe reused have practical value in real systems for economic andenvironmental reasons. However, inexpensive adsorbents, such asindustrial by-products and natural materials, may be used onlyonce and they need not be regenerated for reuse because usually thecosts of desorbing chemicals and regeneration process are higherthan that of these adsorbents.

Desorption of F is carried out by leaching of adsorbed F byacids, bases and salts (Table 3). The selection of a desorbing agentlargely depends on the influence of pH on F adsorption and thestrength of adsorption. Adsorption on most of the adsorbentsdecreases at high pHs, and therefore bases having high pHs arecommonly used for desorption of F from such adsorbents (Table 3).As explained previously, the desorption of F at high pHs is dueto the increased repulsive forces between the negatively chargedadsorbent surface and negatively charged F− ions in solution aswell as the competition between increased concentration of OH−

at high pH values and F− for adsorption. In materials where Fadsorption increases with pH, acids have been found to be moreeffective as desorbing agents. For example, in a hydrous manganeseoxide coated alumina [29] and a thio-urea modified amberlite

resin [34], F adsorption decreased at lower pH values. Consistentwith this adsorption pattern, the percentage of F desorption wasfound to be higher at low pH values. Adsorption of F by mostadsorbents is strong and not easily reversible, partly due to the

1 ardous

caaW[

7

umlitea

belatotaiwap

GaeamAc

aoadaoistow

A

tCe

R

8 P. Loganathan et al. / Journal of Haz

hemical nature of the adsorption process. Therefore, stronger acidsnd bases, as well as a longer time of their interaction with thedsorbent are required for efficient F desorption [36,41,50,55].ater and neutral salts are generally found to be less effective

31,41].

. Conclusions

Adsorption and precipitation/coagulation methods are widelysed for the defluoridation of water. The precipitation/coagulationethod has the drawback of not being efficient in waters having

ow F concentration, and the process produces toxic AlF complexesn the treated water as well as large volumes of waste. The adsorp-ion processes are generally considered attractive because of theirffectiveness, convenience, ease of operation, simplicity of design,nd for economic and environmental reasons.

Multivalent metal oxides and hydroxides and layered dou-le hydroxides generally have high F adsorption capacities. Ionxchange resins and fibres, zeolites, and carbon materials haveow adsorption capacities on their own, but when their surfacesre modified by incorporating organic functional groups or mul-ivalent metallic cations, the adsorption capacity increased. Somef the adsorbents (e.g. layered double hydroxides, alumina) hado be activated by calcining at high temperatures to increase thedsorption capacity. The above materials have potential for usen industrial countries and in areas where the F concentration in

ater is very high. Natural and industrial by-products have lowdsorption capacities but because they are inexpensive, they haveotential for use in rural areas, especially in developing countries.

The pH of water is a dominant factor influencing F adsorption.enerally, F adsorption increases from acidic to near neutral pHnd then decreases with increase in pH. Consistent with this pHffect, acids and alkali have been successfully used to desorb Fnd consequently regenerate the adsorbent for reuse. However,ultiple regeneration and reuse reduce the adsorptive capacity.nother important factor influencing F adsorption is the type andoncentration of other ions present in water.

Future research needs to explore highly efficient, low costdsorbents that can be easily regenerated for reuse over severalperational cycles without significant loss of adsorptive capacitynd have good hydraulic conductivity to prevent filters clogginguring a fixed-bed treatment process. Surface modifications of thedsorbents can be explored to increase the capacity and strengthf adsorption without significantly increasing the cost. The major-ty of studies reported have been conducted in batch trials onynthetic waters. These trials need to be extended using con-inuous mode column trials which have more relevance to realperating systems on natural waters containing other ions asell.

cknowledgements

This study was funded by cooperative Research Centre for Con-amination Assessment and Remediation of the Environment (CRCARE) (Project Number 4.1.12.11/12). We thank Phil Thomas forditing this manuscript.

eferences

[1] J. Fawell, K. Bailey, E. Chilton, E. Dahi, L. Fewtrell, Y. Magara, Fluoride in Drink-ing Water, World Health Organization, IWA Publishing, UK, 2006.

[2] World Health Organisation, Guidelines for Drinking-water Quality, Volume 2.

Health Criteria and Other Supporting Information, second ed., World HealthOrganisation, Geneva, 1996.

[3] U.S. Public Health Service, US Public Health Service Drinking Water Standards,US Government Printing Office, Department of Health Education and Welfare,Washington DC, 1962.

Materials 248– 249 (2013) 1– 19

[4] Y. Cengeloglu, E. Kir, M. Ersöz, Removal of fluoride from aqueous solution byusing red mud, Sep. Purif. Technol. 28 (2002) 81–86.

[5] N.J. Chinoy, Effects of fluoride on physiology of animals and human beings,Ind. J. Environ. Toxicol. 1 (1991) 17–32.

[6] P.T.C. Harrison, Fluoride in water: a UK perspective, J. Fluor. Chem. 126 (2005)1448–1456.

[7] E. Kumar, A. Bhatnagar, J. Minkyu, W. Jung, S. Lee, S. Kim, G. Lee, H. Song, J.Choi, J. Yang, B. Jeon, Defluoridation from aqueous solutions by granular ferrichydroxide (GFH), Water Res. 43 (2009) 490–499.

[8] R.C. Meenakshi, Maheshwari, Fluoride in drinking water and its removal, J.Hazard. Mater. B137 (2006) 456–463.

[9] S. Vigneswaran, C. Visvanathan, Water Treatment Processes. Simple Options,CRC Press, Inc., Florida, 1995.

[10] J.J. Murray, Appropriate Use of Fluorides for Human Health, World HealthOrganisation, Geneva, 1986.

[11] P. Loganathan, M.J. Hedley, N.D. Grace, J. Lee, S.J. Cronin, N.S. Bolan, J.M. Zan-ders, Fertiliser contaminants in New Zealand grazed pasture with specialreference to cadmium and fluorine – a review, Aust. J. Soil Res. 41 (2003)501–532.

[12] Y.H. Li, S. Wang, X. Zhang, J. Wei, C. Xu, Z. Luan, D. Wu, Adsorption of flu-oride from water by aligned carbon nanotubes, Mater. Res. Bull. 38 (2003)469–476.

[13] X. Liao, B. Shi, Adsorption of fluoride on zirconium(IV)-impregnated collagenfiber, Environ. Sci. Technol. 39 (2005) 4628–4632.

[14] S.P. Mishra, M. Das, U.N. Dash, Review on adverse effects of water contami-nants like arsenic, fluoride and phosphate and their remediation, J. Sci. Ind.Res. 69 (2010) 249–253.

[15] E.G. Paulson, Reducing fluoride in industrial wastewater, Chem. Eng. 84(1977) 89–94.

[16] P.D. Nemade, A. Vasudeva Rao, B.J. Alappat, Removal of fluorides from waterusing low cost Adsorbents, Water Science and Technology: Water Supply 2(2002) 311–317.

[17] S. Jagtap, M.K. Yenkie, N. Labhsetwar, S. Rayalu, Fluoride in drinking waterand defluoridation of water, Chem. Rev. (2012) 2454–2466.

[18] W.W. Choi, K.Y. Chen, The removal of fluoride from waters by adsorption, J.Am. Water Works Assoc. 71 (1979) 562–570.

[19] K.B.P.N. Jinadasa, S.W.R. Weerasooriya, C.B. Dissanayake, A rapid method forthe defluoridation of fluoride-rich drinking waters at village level, Inter. J.Environ. Studies. 31 (1988) 305–312.

[20] J.P. Padmasiri, C.B. Dissanayake, A simple defluoridator for removing excessfluorides from fluoride-rich drinking water, Int. J. Environ. Health Res. 5(1995) 153–160.

[21] C. Zevenbergen, L.P. van Reeuwijk, G. Frapporti, R.J. Louws, R.D. Schuiling, Asimple method for defluoridation of drinking water at village level by adsorp-tion on Ando soil in Kenya, Sci. Total Environ. 188 (1996) 225–232.

[22] M.A.M. Sahli, S. Annouar, M. Tahaikt, M. Mountadar, A. Soufiane, A. Elmi-daoui, Fluoride removal for underground brackish water by adsorption onthe natural chitosan and by electrodialysis, Desalination 21 (2007) 37–45.

[23] S. Samatya, U. Yüksel, M. Yüksel, N. Kabay, Removal of fluoride from water bymetal ions (Al3+, La3+ and ZrO2+) loaded natural zeolite, Sep. Sci. Technol. 42(2007) 2033–2047.

[24] K.R. Bulusu, B.B. Sundaresan, B.N. Pathak, W.G. Nawlakhe, D.N. Kulkarni, V.P.Thergaonkar, Fluorides in water, defluoridation methods and their limita-tions, J. Inst. Eng. (Ind.) Environ. Eng. Div. 60 (1979) 1–25.

[25] W.G. Nawlakhe, D.N. Kulkarni, B.N. Pathak, K.R. Bulusu, Defluoridation ofwater with alum by Nalgonda technique, Ind. J. Environ. Health 17 (1975)26–65.

[26] M. Mohapatra, S. Anand, B.K. Mishra, D.E. Giles, P. Singh, Review of fluorideremoval from drinking water, J. Environ. Manage. 91 (2009) 67–77.

[27] A. Bhatnagar, E. Kumar, M. Sillanpää, Fluoride removal from water by adsorp-tion – a review, Chem. Eng. J. 171 (2011) 811–840.

[28] V.K. Gupta, J.M. Carrott, M.M.L.R. Carrott, Suhas, Low-cost adsorbents: Grow-ing approach to wastewater treatment – a review, Crit. Rev. Environ. Sci.Technol. 39 (2009) 783–842.

[29] S. Teng, S. Wang, W. Gong, X. Liu, B. Gao, Removal of fluoride by hydrousmanganese oxide-coated alumina: performance and mechanism, J. Hazard.Mater. 168 (2009) 1004–1011.

[30] S. Meenakshi, N. Viswanathan, Identification of selective ion-exchange resinfor fluoride sorption, J. Colloid Interface Sci. 308 (2007) 438–450.

[31] L. Ruixia, G. Jinlong, T. Hongxiao, Adsorption of fluoride, phosphate, and arse-nate ions on a new type of ion exchange fiber, J. Colloid Interface Sci. 248(2002) 268–274.

[32] F. Helferich, Ion Exchange, Dover Publication Inc., USA, 1995.[33] E.R. Weiner, Applications of Environmental Aquatic Chemistry, CRC Press,

Taylor and Francis Group, Florida, 2008.[34] I.B. Solangi, S. Memon, M.I. Bhanger, An excellent fluoride sorption behaviour

of modified amberlite resin, J. Hazard. Mater. 176 (2010) 186–192.[35] A. Sivasamy, K.P. Singh, D. Mohan, M. Maruthamuthu, Studies on defluorida-

tion of water by coal-based sorbents, J. Chem. Technol. Biotechnol. 76 (2001)717–722.

[36] S.M. Maliyekkal, A.K. Sharma, L. Philip, Manganese-oxide-coated alumina:

a promising sorbent for defluoridation of water, Water Res. 40 (2006)3497–3506.

[37] S. Tokunaga, S.A. Haron, S.A. Wasay, K.F. Wong, K. Laosangthum, A. Uchi-umi, Removal of fluoride ions from aqueous solutions by multivalent metalcompounds, Inter. J. Environ. Studies 48 (1995) 17–28.

ardous

P. Loganathan et al. / Journal of Haz

[38] Y. Wang, E.J. Reardon, Activation and regeneration of a soil sorbent for deflu-oridation of drinking water, Appl. Geochem. 16 (2001) 531–539.

[39] F. Luo, K. Inoue, The removal of fluoride ion by using metal (III)-loaded Amber-lite resins, Solvent Extract. Ion Exch. 22 (2004) 305–322.

[40] M.S. Onyango, Y. Kojima, O. Aoyi, E.C. Bernardo, H. Matsuda, Adsorption equi-librium modelling and solution chemistry dependence of fluoride removalfrom water by trivalent-cation-exchanged zeolite F-9, J. Coloid Interface Sci.279 (2004) 341–350.

[41] M.S. Onyango, Y. Kojima, A. Kumar, D. Kuchar, M. Kubota, H. Matsuda, Uptakeof fluoride by Al3+ pretreated low-silica synthetic zeolites: adsorption equi-librium and rate studies, Sep. Sci. Technol. 41 (2006) 683–704.

[42] R.L. Ramos, J. Ovalle-Turrubiartes, M.A. Sanchez-Castillo, Adsorption of flu-oride from aqueous solution on aluminium-impregnated carbon, Carbon 37(1999) 609–617.

[43] S.A. Wasay, M.J. Haron, S. Tokunaga, Adsorption of fluoride, phosphate, andarsenate ions on lanthanum-impregnated silica gel, Water Environ. Res. 68(1996) 295–300.

[44] Y. Zhou, C. Yu, Y. Shan, Adsorption of fluoride from aqueous solutionon La3+-impregnated cross-linked gelatin, Sep. Purif. Technol. 36 (2004)89–94.

[45] P. Trivedi, L. Axe, Long-term fate of metal contaminants in soils and sedi-ments: role of intraarticle diffusion in hydrous metal oxides, in: R. Hamon, M.McLaughlin, E. Lombi (Eds.), Natural Attenuation of Trace Element Availabilityin Soils, CRC Taylor and Francis Group, New York, 2006, pp. 57–71.

[46] A.K. Chaturvedi, K.P. Yadava, K.C. Pathak, V.N. Singh, Defluoridation of waterby adsorption on fly ash, Water Air Soil Pollut. 49 (1990) 51–61.

[47] S. Ghorai, K.K. Pant, Equilibrium, kinetics and breakthrough studies foradsorption of fluoride on activated alumina, Sep. Purif. Technol. 42 (2005)265–271.

[48] M.G. Sujana, R.S. Thakur, S.B. Rao, Removal of fluoride from aqueous solutionby using alum sludge, J. Colloid Interface Sci. 206 (1998) 94–101.

[49] Y. Ku, H. Chiou, The adsorption of fluoride ion from aqueous solution byactivated alumina, Water Air Soil Pollut. 133 (2002) 349–360.

[50] A.M. Raichur, M.J. Basu, Adsorption of fluoride onto mixed rare earth oxides,Sep. Purif. Tech. 24 (2001) 121–127.

[51] L.E.L. Hammari, A. Laghzizil, P. Barboux, K. Lahlil, A. Saoiabi, Retention of flu-oride ions from aqueous solution using porous hydroxyapatite. Structure andconductive properties, J. Hazard. Mater. B114 (2004) 41–44.

[52] X. Fan, D.J. Parker, M.D. Smith, Adsorption kinetics of fluoride on low costmaterials, Water Res. 37 (2003) 4929–4937.

[53] H. Wang, J. Chen, Y. Cai, J. Ji, L. Liu, H.H. Teng, Defluoridation of drinking waterby Mg/Al hydrotalcite-like compounds and their calcined products, Appl. ClaySci. 35 (2007) 59–66.

[54] L. Lv, J. He, M. Wei, D.G. Evans, X. Duan, Factors influencing the removal of fluo-ride from aqueous solution by calcined Mg-Al-CO3 layered double hydroxides,J. Hazard. Mater. B133 (2006) 119–128.

[55] D.P. Das, J. Das, K. Parida, Physicochemical characterization and adsorp-tion behaviour of calcined Zn/Al hydrotalcite-like compound (HTlc) towardsremoval of fluoride from aqueous solution, J. Colloid Interface Sci. 261 (2003)213–220.

[56] C. Diaz-Nava, M. Solache-Rios, M.T. Olguin, Sorption of fluoride ions fromaqueous solutions and well drinking water by thermally treated hydrotalcite,Sep. Sci. Technol. 38 (2003) 131–147.

[57] Y. Ku, H. Chiou, W. Wang, The removal of fluoride ion from aqueous solutionby a cation synthetic resin, Sep. Sci. Technol. 37 (2002) 89–103.

[58] K.M. Popat, P.S. Anand, B.D. Dasare, Selective removal of fluoride ions fromwater by the aluminium form of the aminomethylphosphonic acid-type ionexchanger, React. Polym. 23 (1994) 23–32.

[59] A.A.M. Daifullah, S.M. Yakout, S.A. Elreefy, Adsorption of fluoride in aque-ous solutions using KMnO4-modified activated carbon derived from steampyrolysis of rice straw, J. Hazard. Mater. 147 (2007) 633–643.

[60] I. Abe, S. Iwasaki, T. Tokimoto, N. Kawasaki, T. Nakamura, S. Tanada, Adsorp-tion of fluoride ions onto carbonaceous materials, J. Colloid Interface Sci. 275(2004) 35–39.

[61] Y.H. Li, S. Wang, X. Zhang, J. Wei, C. Xu, Z. Luan, D. Wu, B. Wei, Removal offluoride from water by carbon nanotube supported alumina, Environ. Technol.24 (2003) 391–398.

[62] P.D. Nemade, A.V. Rao, B.J. Alappat, Removal of fluorides from water usinglow cost adsorbents, Water Sci. Tech. Water Supply 2 (2002) 311–317.

[63] N. Viswanathan, S. Meenakshi, Selective fluoride adsorption by a hydrotal-cite/chitosan composite, Appl. Clay Sci. 48 (2010) 607–611.

[64] C.S. Sundaram, N. Viswanathan, S. Meenakshi, Fluoride sorption by nano-hydroxyapatite/chitin composite, J. Hazard Mater. 172 (2009) 147–151.

[65] C.S. Sundaram, N. Viswanathan, S. Meenakshi, Uptake of fluoride by nano-hydroxyapatite/chitosan, a bioinorganic composite, Bioresource Technol. 99(2008) 8226–8230.

[66] W. Ma, F. Ya, M. Han, R. Wang, Characteristics of equilibrium, kinetics studies

for adsorption of fluoride on magnetic-chitosan particle, J. Hazard. Mater. 143(2007) 296–302.

[67] S.P. Kamble, S. Jagtap, N.K. Labhsetwar, D. Thakare, S. Godfrey, S. Devotta,S.S. Rayalu, Defluoridation of drinking water using chitin, chitonsan andlanthanum-modified chitosan, Chem. Eng. J. 129 (2007) 173–180.

Materials 248– 249 (2013) 1– 19 19

[68] M.G. Sujana, H.K. Pradhan, S. Anand, Studies on sorption of some geomaterialsfor fluoride removal from aqueous solutions, J. Hazard. Mater. 161 (2009)120–125.

[69] M. Agarwal, K. Rai, R. Shrivastav, S. Dass, Defluoridation of water usingamended clay, J. Cleaner Prod. 11 (2003) 439–444.

[70] M. Srimurali, A. Pragathi, J. Karthikeyan, A study on removal of fluorides fromdrinking water by adsorption onto low-cost materials, Environ. Pollut. 99(1998) 285–289.

[71] M. Islam, R. Patel, Thermal activation of basic oxygen furnace slag and evalu-ation of its fluoride removal efficiency, Chem. Eng. J. 169 (2011) 68–77.

[72] A. Tor, N. Danaoglu, G. Arslan, Y. Cengeloglu, Removal of fluoride from waterby using granular red mud: batch and column studies, J. Hazard. Mater. 164(2009) 271–278.

[73] V.K. Gupta, I. Ali, V.K. Saini, Defluoridation of wastewaters using waste carbonslurry, Water Res. 41 (2007) 3307–3316.

[74] R. Piekos, S. Paslawska, Fluoride uptake characteristics of fly ash, Fluoride 32(1999) 14–19.

[75] Y.D. Lai, J.C. Liu, Fluoride removal from water with spent catalyst, Sep. Sci.Technol. 31 (1996) 2791–2803.

[76] Y. Tang, X. Guanc, J. Wang, N. Gaob, M.R. Mcphaild, C.C. Chusuei, Fluorideadsorption onto granular ferric hydroxide: effects of ionic strength, pH, sur-face loading, and major co-existing anions, J. Hazard. Mater. 171 (2009)774–779.

[77] Y. Zhou, R.J. Haynes, Sorption of heavy metals by inorganic and organic com-ponents of solid wastes: Significance to use of wastes as low-cost adsorbentsand immobilising agents, Crit. Rev. Environ. Sci. Technol. 40 (2011) 909–977.

[78] N. Chubar, Physico-chemical treatment of micropollutants: adsorption andion exchange, in: J. Virkutyte, R.S. Varma, V. Jegatheesan (Eds.), Treatmentof Micropollutants in Water and Wastewater, IWA publishing, London, 2010,pp. 165–203.

[79] W.H. Zou, R.P. Han, Z.Z. Chen, J. Shi, H.M. Liu, Characterization and propertiesof manganese oxide coated zeolite as adsorbent for removal of copper (II) andlead (II) ions from solution, J. Chem. Eng. Data 51 (2006) 534–541.

[80] T. Takamatsu, M. Kawashima, M. Koyama, The role of Mn2+ rich manganeseoxide in the accumulation of arsenic in lake sediments, Water Res. 9 (1985)1029–1032.

[81] E.N. Peleka, E.A. Deliyanni, Adsorptive removal of phosphates from aqueoussolutions, Desalination 245 (2009) 357–371.

[82] K. Goh, T. Lim, Z. Dong, Application of layered double hydroxides for removalof oxyanions: a review, Water Res. 42 (2008) 1343–1368.

[83] K. Vaaramma, J. Lehto, Removal of metals and anions from drinking water byion exchange, Desalination 155 (2003) 157–170.

[85] C. Diaz-Nava, M.T. Olguin, M. Solache-Rios, Water defluoridation by Mexicanheulandite-clinoptilolite, Sep. Sci. Technol. 37 (2002) 3109–3128.

[86] P. Chingombe, B. Saha, R.J. Wakeman, Surface modification and characterisa-tion of a coal-based activated carbon, Carbon 43 (2005) 3132–3143.

[87] C.Y. Yin, M.K. Aroua, W.M.A.W. Daud, Review of modifications of activatedcarbon for enhancing contaminant uptakes from aqueous solutions, Sep. Purif.Technol. 52 (2007) 403–415.

[88] R. Mahmudov, C.P. Huang, Perchlorate removal by activated carbon adsorp-tion, Sep. Purif. Technol. 70 (2010) 329–337.

[89] A. Bhatnagar, M. Sillanpää, A review of emerging adsorbents for nitrateremoval from water, Chem. Eng. J. 168 (2011) 493–504.

[90] C.S. Sundaram, N. Viswanathan, S. Meenakshi, Defluoridation chemistry ofsynthetic hydroxyapatite at nano scale: equilibrium and kinetic studies, J.Hazard Mater. 155 (2008) 206–215.

[91] A. Bhatnagar, V.J.R. Vilar, C.M.S. Botelho, R.A.R. Boaventura, A review of theuse of red mud as adsorbent for the removal of toxic pollutants from waterand wastewater, Environ. Technol. 32 (2011) 231–249.

[92] D. Mohan, C.U. Pittman Jr., Arsenic removal from water/wastewater usingadsorbents – a critical review, J. Hazard. Mater. 142 (2007) 1–53.

[93] H. Genc-Fuhrman, H. Bregnhoj, D. McConchie, Arsenate removal from waterusing sand-red mud columns, Water Res. 39 (2005) 2944–2954.

[94] C. Zhu, Z. Luan, Y. Wang, X. Shan, Removal of cadmium from aqueous solu-tions by adsorption on granular red mud (GRM), Sep. Purif. Technol. 57 (2007)161–169.

[95] I.A.M. Yunusa, P. Loganathan, S.P. Nissanka, V. Manoharan, M.D. Burchett,C.G. Skilbert, D. Eamus, Application of coal fly ash in agriculture: a strategicperspective, Crit. Rev. Environ. Sci. Technol. 42 (2012) 559–600.

[96] W. Nigussie, F. Zewge, B.S. Chandravanshi, Removal of excess fluoride fromwater using waste residue from alum manufacturing process, J. Hazard. Mater.147 (2007) 954–963.

[97] R.H. Masue, T.A. Loeppert, T.A. Kramer, Arsenate arsenite adsorption anddesorption behaviour on coprecipitated aluminium:iron hydroxides, Environ.Sci. Technol. 41 (2007) 837–842.

[98] L. Yan, Y. Xu, H. Yu, X. Xin, Q. Wei, B. Du, Adsorption of phosphate fromaqueous solution by hydroxyl-aluminum, hydroxyl-iron and hydroxyl-iron-

aluminium pillared bentonites, J. Hazard. Mater. 179 (2010) 244–250.

[99] S.D. Faust, O.M. Aly, Adsorption Processes for Water Treatment, Butterworths,Boston, 1989.

[100] N.Y. Mezenner, A. Bensmaili, Kinetics and thermodynamic study of phosphateadsorption on iron hydroxide-eggshell waste, Chem. Eng. J. 147 (2009) 87–96.