Adsorption of Co Ni Cu and Zn on Amorphous Hydrous MnO2 From 1-1 Electrplyte Solutions-libre

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Journal of Colloid and Interface Science 269 (2004) 11–21

www.elsevier.com/locate/jcis

Adsorption of Co2+, Ni2+, Cu2+, and Zn2+ onto amorphous hydrousmanganese dioxide from simple (1–1) electrolyte solutions

Sukriti Bhusan Kanungo,1 Sushree Swarupa Tripathy,∗ Santosh Kumar Mishra,Biswanath Sahoo, and Rajeev

Regional Research Laboratory, Bhubaneswar 751013, Orissa, India

Received 17 March 2003; accepted 24 July 2003

Abstract

The adsorption of Co2+, Ni2+, Cu2+, and Zn2+ onto amorphous hydrous manganese dioxide (δ-MnO2) has been studied using two

methods, viz., isotherms at constant pH in the presence of buffer solution and pH variation in the absence of buffer solution from a fixed

metal ion concentration. While the adsorption isotherm experiments were carried out in 0.5 M NaCl only, pH variation or batch titration

experiments were carried out in 0.5 M NaCl, 0.01 M NaCl, and 0.01 M KNO3 solutions. The complex nature of adsorption isotherms at

constant pH values indicates that adsorption of all the cations is non-Langmuirian (Freundlich) and takes place on the highly heterogeneous

oxide surface with different binding energies. The proton stoichiometry derived from isotherms at two close pH values varies between 0.3

and 0.8. The variation of fractional adsorption with pH indicates that the background electrolyte solution influences the adsorption of cations

through either metal-like or ligand-like complexes with Cl−, the former showing a low adsorption tendency. The proton stoichiometry values

derived from the Kurbatov-type plot varies not only with the electrolyte solution but also with the adsorbate/adsorbent ratio. The variation

of fractional adsorption with pH can be modeled either with the formation of the SOM + type or with a combination of SOM+ and SOMOH

type complexes, depending upon the cation and electrolyte medium. The equilibrium constants obtained from Kurbatov-type plots are found

to be most suitable in these model calculations. Adsorption calculated on the basis of ternary surface metal–chlorocomplex formation exhibits

very low values.

© 2003 Elsevier Inc. All rights reserved.

Keywords: Adsorption; Metal ions; δ-MnO2; Simple electrolyte medium

1. Introduction

Adsorption of trace metals on hydrous manganese diox-

ide is a widely studied topic of research from the point of

view of environmental and geochemical aspects. Though the

physicochemical phenomena involved in both the aspects arethe same, the main difference is the period for which the ad-

sorbate and adsorbent are in the state of equilibrium. While

the environmental chemists are more interested in shorter

periods, geochemical processes generally involve long-term

interaction or dynamic states of equilibrium, where either

adsorbent or adsorbate is in continuous supply in the nat-

ural environment [1,2]. However, the distinction is not rigid

* Corresponding author.

E-mail address: [email protected] (S.S. Tripathy).1 Present address: Flat No. A7/2, Konnagar Abasan, Konnagar 712235,

District Hooghly, West Bengal, India.

and environmental chemistry also deals with long-term geo-

chemical interaction.

Despite many regulatory measures, various industrial

waste products, both inorganic and organic, are being dis-

charged into natural water systems, mainly because of eco-

nomic compulsions. To understand either the self-cleaningor the induced cleaning capacity of particulate matters in

natural water systems we need to know the phenomeno-

logical behavior of solute–adsorbent interaction in such

systems. Basically, the same approach is also applied to un-

derstand the geochemical processes at the mineral surface,

but the system is more complex because of the surrounding

conditions such as Eh, pH, atmospheric oxygen, and car-

bon dioxide, which play important roles in controlling the

processes [1–3].

Amorphous δ-MnO2, either detrital or authigenic, is one

of the important constituents of particulate matter in many

natural water systems, especially in sea water. Because of

0021-9797/$ – see front matter © 2003 Elsevier Inc. All rights reserved.

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12 S.B. Kanungo et al. / Journal of Colloid and Interface Science 269 (2004) 11–21

its high surface charge, hydrous MnO2 has high scavenging

power for trace metals. It has been generally observed that

transition metal ions adsorb onto this oxide mainly through

specific chemical interaction, even by releasing Mn2+ in

solution from its disordered lattice [4–8]. This shows that

unlike other oxide substrates adsorption of cations onto hy-

drous MnO2 is reasonably complicated and therefore nosimple proton stoichiometry is likely to follow over a wide

range of pH. The situation is further complicated if hydrous

MnO2 is associated with other oxide substrates such as iron

oxide, aluminum oxide, silica, clay minerals etc. as found in

natural environments.

In the present paper an attempt has been made to study

the sorption of Co2+, Ni2+, Cu2+, and Zn2+ onto hydrous

MnO2 (δ-MnO2) in simple 1–1 electrolyte media with a

view to examining (i) adsorption isotherms (increasing ad-

sorbate concentration) at constant pH values, (ii) the nature

of solute–adsorbent interaction in the presence of different

electrolyte media, and (iii) the extent to which the surfacecomplexation models that form the basis of our present un-

derstanding of the adsorption of inorganics are valid for

the present system involving hydrous manganese dioxide,

whose surface is generally more heterogeneous than many

common adsorbents like, iron oxide or aluminum oxide.

2. Experimental

2.1. Materials

The method of preparation of the δ-MnO2 sample and itscharacterization by different physicochemical methods have

been described in our previous communication [9]. All the

chemicals used were of analytical grade. All the aqueous so-

lutions were prepared in all-glass double distilled water.

2.2. Adsorption procedure (batch titration in presence of

metal ions)

To a series of thoroughly cleaned and dried 125-ml

polypropylene bottles, 45 ml of electrolyte solution contain-

ing a known quantity of metal ion was transferred. The pH

values of the solutions were first adjusted from 2 to 8 byadding known volumes of either 0.1 M HCl or 0.1 M NaOH.

The total volume of solution in each bottle was then made up

to 50 ml by adding the relevant electrolyte solution. The pH

values of the solutions were then noted accurately. A known

amount (0.01–0.02 g) of adsorbent sample was then added

to each bottle. The suspensions were bubbled with flowing

nitrogen gas (99%) at the flow rate of about 50–60 ml per

minute for about 5 min and immediately capped, shaken

well, and left to equilibrate at room temperature (300 K),

during which the bottles were shaken intermittently. After

72 h of equilibration about 15 ml of clear supernatant liq-

uid was carefully decanted off and centrifuged for 15 min.

Metal contents including Mn were estimated in the super-

natant liquid by atomic absorption spectrophotometry (Var-

ian AA+) using air–acetylene flame. The final pH value of

the equilibrated solution in each bottle was determined with

the help of EDT (England) pH/ion meter. Standard metal ion

solutions were prepared in the same background electrolyte

solution in which batch titration was carried out. A blank titration was also carried out in the similar manner without

adding any metal ion.

2.3. Adsorption isotherm

Adsorption experiments with increasing concentrations

of metal ion at a constant temperature (300 K) and pH value

were carried out in the same manner as described above, ex-

cept that instead of a fixed quantity of metal ion solution in

each bottle, varying quantities of metal ion solutions were

added. Initially, an attempt was made to maintain constant

pH by adding either 0.1 M NaOH or 0.1 M HCl. However,this procedure was not successful in maintaining the desired

pH value, which always tended to drop gradually, especially

at high metal ion concentrations even after the addition of

large quantities of alkali solution. Therefore, all the isotherm

experiments were carried out in 50 ml of universal buffer so-

lutions at different pH values [10]. Because of the special

nature of the work, the buffer solutions were prepared ac-

cording to slightly modified procedure as follows.

2.4. Solution of mixed acids (solution A)

A sample of 4.6 ml of 17.4 N CH3COOH, 5.84 ml of

41.1 N H3PO4, and 4.95 g of H3BO3 (M.W. 61.83 g) was

taken in a 2-L volumetric flask into which 1 L of 1 N NaCl

solution was added. The solution was thoroughly shaken un-

til clear solution was obtained. The total volume was made

up to 2 L with distilled water.

2.5. Solution with 0.2 M NaOH (solution B)

A sample of 16 g of sodium hydroxide pellets was dis-

solved in 0.5 M NaCl and after cooling diluted to 2 L in a

volumetric flask with 0.5 M NaCl solution. The solutions A

and B were mixed as follows to make buffer solution of dif-

ferent pH values:

pH Volume of Volume of Total

solution A solution B volume

(ml) (ml) (ml)

4.04 600 165 765

4.49 600 210 810

5.05 600 255 855

5.60 600 270 870

6.05 600 300 900

7.00 600 390 990

The standard solutions of trace metal ions were also pre-

pared in the corresponding buffer solutions. The initial and

final pH values remain almost the same (±0.5 unit) in the



S.B. Kanungo et al. / Journal of Colloid and Interface Science 269 (2004) 11–21 13

presence of buffer solution even at the highest concentration

of metal.

2.6. Data accuracy

All the adsorption data were obtained with respect to

the original concentration of metal (control) prepared in the

same medium used in the relevant adsorption experiment.

This procedure of using control solution was followed for

every experimental determination of metal by AAS. This

reduces considerably the absolute error associated with the

AAS method from ±8.0% to ±5.0%. This relative error is

further reduced (< ±3.0%) when adsorption is expressed in

relative term such as percent adsorption.

3. Results and discussion

3.1. Adsorption isotherm

Figures 1–4 illustrate the adsorption isotherms of Co2+,

Ni2+, Cu2+, and Zn2+, respectively, in 0.5 M NaCl solu-

tion and at different pH values. The high concentration of

NaCl was selected primarily because of our ultimate objec-

tive of studying adsorption in sea water. The figures show

that adsorption of cation increases with increased pH and for

Cu2+ and Zn2+ the adsorption at pH 6.0 was so high even

at low concentrations that a separate scale became necessary

to represent the data (see Figs. 3 and 4). The notable feature

in the isotherms of all the four cations is the appearance of

more than one step, indicating adsorption at sites with differ-

ent binding energies. In the case of Co2+ the concentration

at which the first break occurs increases from 0.05 mM at

pH 4.45 to 0.15 mM at pH 5.05. For these two pH values,

adsorption continues to increase with increased equilibrium

concentration without the appearance of distinct second sat-

uration plateau. The adsorption isotherm of Ni2+ has close

similarity to that of Co2+ except that in the case of Ni2+

two distinct steps of saturation plateau appear at pH 4.13.

For both the cations the first break in the isotherms at pH

6.05 appears in the form of a minor shoulder followed by

gradual attainment of a broad saturation region. In contrast,

the first plateau region in the adsorption isotherms of Cu 2+

and Zn2+ is followed by a sharp rise in adsorption. Interest-

ingly, for the adsorption in Cu2+ the first saturation plateau

is spread over a very wide concentration range unlike any

other cations.

As far as the role of buffer solution on the adsorption

of cations is concerned, it is difficult to draw any definite

conclusion in the absence of specific data for the adsorp-

tion of anions of buffer solution which is beyond the scope

of the present work. As the pH values of the isotherms are

well above pHpzc of δ-MnO2 (pHpzc 1.8), any significant ad-

sorption of anion is ruled out. However, by comparing the

adsorption data obtained from the batch titration at different

Fig. 1. Adsorption isotherm of Co2+ in 0.5 M NaCl solution at different pH

values.

Fig. 2. Adsorption isotherm of Ni2+ in 0.5 M NaCl solution at different pH

values.

Fig. 3. Adsorption isotherm of Cu2+ in 0.5 M NaCl solution at different pH

values.




Fig. 4. Adsorption isotherm of Zn2+ in 0.5 M NaCl solution at different pH

values.

pH values, i.e., absence of buffer solution, with those ob-

tained from adsorption isotherms in the presence of buffer

solution at the same pH values and initial concentration of

metal ions (determined separately in duplicate), it is possi-

ble to examine the role of the buffer ions on adsorption. The

relevant data are summarized in Table 1 from which the fol-

lowing observations can be made.

(i) At pH 6.05 there is little difference in the adsorption of

metal ions in presence and in absence of buffer solu-

tion. As the pH values decreases the adsorption in the

buffered medium tends to decrease more than in the un-

buffered medium. At pH 4.0 the difference is highest

for all the four metal ions. This suggests that at low pH

all the metal ions tend to form complexes with anions

such as acetate and phosphate in solution phase, thereby

reducing the adsorption capacity of the metal ions.

(ii) In the case of Zn2+ the decrease in adsorption in the

buffered medium is highest at low pH indicating that

Zn2+ forms more stable complexes with buffer ions.

(iii) While Ni2+ shows small variation in adsorption with

the nature of the medium, Cu2+ appears to show more

adsorption in the buffered medium. It is possible that

Ni2+ forms more stable complexes with Cl− than withbuffer anions and therefore the later have little effect.

On the other hand, in addition to hydroxy species Cu2+

may form complex species with buffer ions, especially,

acetate ions which are adsorbed in a metal-like manner

with ligand ions exposed to the solution phase [7]. This

creates more adsorption sites and stimulates further ad-

sorption of Cu2+ [11].

Though the stability constant values of the metal ions in Ta-

ble 2 broadly follows Irving–William order (Co2+ < Ni2+ <

Cu2+ > Zn2+) the adsorption behavior follows the order

Ni2+ < Zn2+ < Co2+ < Cu2+ (see Table 1). Similar se-quence of adsorption order has also been observed by other

workers [16,17] on hydrous MnO2. This deviation from the

Irving–William order is mainly for two reasons: (a) Surface

complexation does not occur through purely chemical in-

teraction like any common complexing ligand. Electrostatic

force plays also an important role. (b) Cu2+ forms several

hydrolytic and polymeric species in solution which greatly

enhances adsorption onto hydrous MnO2 surface. Such a

phenomenon is not observed for the other three metal ions,

at least within the pH range 3.0–7.0.

An attempt has been made to linearize the isotherm by

plotting adsorption (mol/l) vs concentration (mol/l) on log–log scale. The same unit has been used for both the axes

with the objective of deriving the proton stoichiometry, as

the amount of adsorbent used for a particular cation is same

for all the pH values. An example of such plot is illustrated

in Fig. 5 which shows that the slope values in the initial stage

of low concentration are much lower than unity, indicating

the non-Langmuirian nature of adsorption. For the adsorp-

tion isothermsof Ni2+, Cu2+, and Zn2+, the log–log plots of

adsorption vs concentration are not shown because of space

constraints. However, the slope values of the initial linear

regions of these plots are much lower than unity indicating

that adsorption is not proportional to concentration. From

the initial linear regions at two close pH values, proton sto-

Table 1

Comparison of the extent of adsorption obtained from isotherms at different pH values with those from batch titration in the presence of trace metal ions

Average Co2+ (1.620× 10−4 M) Ni2+ (1.704× 10−4 M) Cu2+ (1.500× 10−4 M ) Zn2+ (1.530× 10−4 M )

pH value Adsorption Batch Adsorption Batch Adsorption Batch Adsorption Batch

isotherm titration isotherm titration isotherm titration isotherm titration

mmol/g mmol/g mmol/g mmol/g mmol/g mmol/g mmol/g mmol/g

6.05 0.7375 0.7371 0.3608 0.3876 0.7600 0.7275 0.6650 0.6962

5.60 – – – – – – 0.5085 0.6525

5.05 0.6387 0.6725 0.3425 0.3365 0.7411 0.7050 0.4153 0.5815

4.45 0.4661 0.5570 0.3308 0.3067 0.5860 0.6000 0.3200 0.4781

4.13 – – 0.2573 0.2770 – – – –




Table 2

Stability constants (log Q) values of the complexes formed by the relevant trace metal ions with some buffer anions in NaCl solution of different ionic strengths

Metal ion MAC+ MHPO4 MCl+ MOH+

(M2+) I = 0a I = 0.67 Mb I = 0.1 Ma I = 0c I = 0.67 Md I = 0e I = 0.67 Md

Co2+ −3.0 −1.80 −2.18 −0.57 −0.08 −4.35 −4.00

Ni2+ −1.12 −1.74 −2.08 −0.72 0.0 −4.14 −3.80

Cu2+ −2.24 −2.40 −2.40 −0.40 −0.34 −6.50 −5.56

Zn2+ −1.57 −2.00 −2.40 −0.49 −0.29 −5.04 −4.80

Note. I = 0.67 M (sea water).a J.A. Dean, Langes Handbook of Chemistry, 13th ed., McGraw Hill, 1987, ch. 5, p. 71.b L. Balistrieri, P.G. Bower, J.W. Murray, Deep-Sea Res. A 28 (1981) 101.c D.L. Turner, M. Whitefield, A.G. Dickson, Geochim. Cosmochim. Acta 45 (1981) 855.d H. Rupert, Chem. Erde 39 (1980) 97.e C.F. Baes, R.E. Mesmer, The Hydrolysis of Cations, Wiley, New York, 1976. Quoted in D.A. Dzombak, F.M.M. Morel, Surface Complexation Modeling.

Hydrous Iron Oxide, Wiley, New York, 1990, pp. 104–105.

Fig. 5. Log–log plot of equilibrium concentration of Co2+ and its adsorp-

tion (mol/l) from 0.5 M NaCl solution.

ichiometry can be estimated by using the method proposed

by Perona and Leckie [12]:

(1)

∂Γ H+

∂ Γ M2+

pH

=

∂ log[M2+]soln

∂pH

Γ

M2+

= χp.

The left-hand side of the above equation represents the

change in H+ adsorbed with the metal ion adsorbed at

constant pH, i.e., proton stoichiometry (χp). This can be

obtained from the change in metal ion concentration at aconstant adsorption density from the initial linear region of

log–log plot of adsorption vs concentration for two close

pH values. If the plots at two pH values are not parallel

to each other, the χp value will vary with adsorption den-

sity. For Co2+ the three isotherms run parallel to each other

(see Fig. 5), both before and after the break. At low con-

centrations, i.e., before the break, χp is 0.2, whereas after

the break, the value varies between 0.36 and 0.5. In the

case of Ni2+, χp values obtained from the linear isotherm

(Freundlich) at pH 7.0, 6.05, and 5.05 at low equilibrium

concentration vary between 0.22 and 0.27. At higher con-

centrations, i.e., after the break, χp value varies widely from

0.32 to 0.60. From the isotherm of Cu2+ at pH 5.1 and

4.45, χp in the initial stage is about 0.6, which varies at

higher concentration. In the case of Zn2+, χp value esti-

mated from the linearized isotherms at pH 4.5 and 5.1 variesfrom 0.80 to 0.97 with increased concentration. Except for

Zn2+, the proton stoichiometry values obtained from the lin-

earized isotherm (Freundlich) by the above method are close

to those obtained from the Kurbatov-type plot as discussed

in a later section of this paper.

The proton stoichiometry results indicate that the adsorp-

tion of trace metal ions on δ-MnO2 does not follow any

stoichiometric reaction involving the release of equivalent

H+ irrespective of the nature of electrolyte medium. This

suggests that buffer anions do not exert any significant in-

fluence in the presence of strong NaCl solution. Table 2

gives some stability constants (log Q) data for acetate andphosphate (HPO2−

4 ) complexes besides chloro and hydroxy

complexes of Co2+, Ni2+, Cu2+, and Zn2+ ions.

The data in Table 2 suggest that monochloro complexes

have highest stability in strong NaCl media such as sea wa-

ter. These are followed by acetato complexes, the stability

constants of which in sea water medium are calculated from

the following relationship found by Balistrieri et al. [13]:

(2)log ∗KMAC = 0.27log∗KMOH − 0.62,

where log ∗KMAC = constant for metal–acetato complex in

solution; log ∗KMOH = first hydrolysis constant of metal in

sea water medium.The ∗KMOH values are taken from column 7 in Table 2.

We have considered the formation of MHPO4, as HPO2−4 is

the most stable species in the pH range 3–7. The data for

the complexation constants of boric acid are not easily avail-

able, but they are certainly much lower. It may, therefore, be

concluded that buffer ions have no major influence on the

adsorption of cations in the presence of strong NaCl solu-

tion.

An attempt has been made to fit the isotherm to a pH-

dependent Langmuir type equation [14,15] as follows,

(3)

[M2+]soln

[M2+]ad=

[H+]

∗K1Γ max+[M2+]

Γ max ,




Table 3

Parameters derived from Langmuir-type plots of the adsorption isotherms of Co 2+ , Ni2+ , Cu2+ , and Zn2+ in 0.5 M NaCl at different pH values

pH Co2+ Ni2+ Cu2+ Zn2+

(Av) Concentration − log∗K1 Γ max Concentration − log∗K1 Γ max Concentration − log∗K1 Γ max Concentration − log ∗K1 Γ max

range mmol/g range mmol/g range mmol/g range mmol/g

(mM) (mM) (mM) (mM)

6.05 0–0.75 4.93 2.50 0–0.20 4.20 0.588 0–0.013 3.69 0.769 0–0.023 4.16 1.087

5.60 – – – – – – – – – 0–0.125 3.72 0.667

5.05 0–0.26 3.79 1.43 0–0.25 3.43 0.500 0–0.012 2.88 0.870 0–0.10 3.02 0.426

4.45 0–0.16 3.0 0.286 – – – 0–0.7 2.52 0.385 0–0.150 2.65 0.347

4.13 – – – 0–0.175 2.44 0.312 – – – – – –

− log∗K1 value

extrapolated to pHpzc 0.60 0.82 0.95 0.80

where ∗K1 = binding constant for adsorption of M2+ as

SOM+; Γ max = maximum adsorption at the relevant pH.

Since the non-Langmuirian behavior is generally attributed

to multisite adsorption, ∗K1 is an average binding constant,

which also involves the release of a single proton from the

surface. The results in Table 3 indicate that while for the ad-sorption of Co2+, Ni2+, and Zn2+, Γ max increases with pH

increase from 4.13 to 6.05, for the adsorption of Cu 2+ an

opposite behavior is observed. This is possibly due to the

fact that hydrolytic species tend to adsorb on the surface of

δ-MnO2 leading to gradual neutralization of surface charge

(SOMOH) and therefore further increase in pH has little ef-

fect on the adsorption density.

Equation (3) also provides us with the average surface

complexation constant (∗K1) from the slope and intercept

values of the initial linear region, i.e., at low concentrations.

The data in Table 3 show that ∗K1 value increases with de-

creased pH. But they tend to converge into a single value forall the cations at pH below 4.0 and on further extrapolation

to pHpzc (pH 1.8) by smooth curves they give − log ∗K int1

values varying between 0.7 and 1.0 for different cations (cf.

Fig. 6). These values are close to those obtained from Kur-

batov plots, which will be discussed in a later section.

3.2. Effect of pH vis-à-vis electrolyte medium on the

adsorption of cations

The most important single factor controlling the adsorp-

tion of metal ions onto hydrous oxides is the pH of the

medium, as evident from the results in the preceding section.

It is generally observed that adsorption of metal ion onto anoxide surface increases sharply and reaching the maximum

value of 100% within a narrow pH range of 1–2 pH units.

This is known as “adsorption edge.” But in the present sys-

tem with δ-MnO2 as adsorbent the adsorption edge is not so

steep, particularly in 0.5 M NaCl. Similar broad adsorption

edges for Co2+, Ni2+, Cu2+, and Zn2+ adsorption onto δ-

MnO2 has also been observed by Murray [6]. Figures 7–9

indicate that adsorption of Co2+, Ni2+, and Cu2+ is highest

in the presence of 0.01 M NaCl and lowest in 0.5 M NaCl.

Adsorption in 0.01 M KNO3 lies intermediate between the

two electrolyte media. In the case of Zn2+ (Fig. 10), adsorp-

tion is highest in 0.5 M NaCl at pH less than 3.5 and lowest

Fig. 6. Plot of apparent binding constants for different cations derived from

Langmuir-type plots (Table 3) as function of pH. The curves when extrapo-

lated to pHpzc give the intrinsic binding constants.

in 0.01 M KNO3, with adsorption in 0.01 M NaCl lying in-

termediate between the two electrolyte solutions. However,

at pH above 5.0 fractional adsorption in all the three elec-

trolyte media almost merge together. This clearly suggests

that Cl− enhances the adsorption of Zn2+ on δ -MnO2.

It has been stated earlier that low proton stoichiometry of

the adsorption of heavy metal ions on to δ -MnO2 indicates

the possible occurrence of some adsorption reactions with-

out the release of proton, irrespective of ionic strength and

the nature of electrolyte (1–1) solution. This suggests thatit is the high surface charge on hydrous MnO2 that leads

to the adsorption not only as free metal ion, but also as

some of its anion complexes. In the previous paper [9] it has

been shown that the concentration of free negatively charged

surface species (SO−) increases sharply while the concen-

tration of ion pairs (SO−Na+) decreases considerably with

decreased ionic strength of the medium. The former tends

to adsorb not only free metal ions but also its complexated

species with H+ still continuing to be associated as charge

balancing counterion. At high electrolyte concentrations ad-

sorption takes place in the same manner as above, but releas-

ing Na+

in the bulk solution.




Fig. 7. Variation of fractional adsorption of Co2+ from different elec-

trolyte solutions as function of pH. Solid line represents model fit

for adsorption in 0.5 M NaCl ([Co2+]t = 1.62 × 10−4 M) using

p∗K1 = 0.9. Dashed line represents model fit for adsorption in 0.01 M

NaCl ([Co2+]t = 2.02×10−4 M) using a combination of p∗K1 = 0.72 and

p∗K4 = 3.0. Dot–dashed line represents model fit for adsorption in 0.01 M

KNO3 ([Co2+]t = 1.92× 10−4 M) using p∗K1 = 0.65.

Fig. 8. Variation of fractional adsorption of Ni2+ from different elec-

trolyte solutions as function of pH. Solid line represents model fit

for adsorption in 0.5 M NaCl ([Ni2+]t = 1.86 × 10−4 M) using

p∗K1 = 0.76. Dashed line represents model fit for adsorption in 0.01 M

NaCl ([Ni2+]t = 1.88× 10−4 M) using a combination of p∗K1 = 1.0 and

p∗K4 = 3.3. Dot–dashed line represents model fit for adsorption in 0.01 M

KNO3 ([Ni2+]t = 1.82 × 10−4 M) using p∗K1 = 1.6 and p∗K4 = 3.65.

Dotted line represents model fit for adsorption in 0.5 M NaCl using

p∗KintML(1)

=−0.1.

The effect of anion (including buffered medium) on theadsorption of cations manifests itself in two ways, viz., ad-

sorption of metal-like (type I) and ligand-like (type II) com-

plexes from the solution phase [11,16] as illustrated below:

Type I

(4)SOH0 +M2++ L−→ SOML0 +H+,

(5)SOH0 +M2++ L−→ SOHML+,

(6)SOH0 +ML02 → SOHML++ L−.

Type II

(7)SOH0

+ L−

+M2+

→ SLM+

+OH−

.

Fig. 9. Variation of fractional adsorption of Cu2+ from different electrolyte

solutions as function of pH. Solid line represents model fit for adsorption in

0.5 M NaCl ([Cu2+]t = 1.50×10−4 M) with a combination of p∗K1 = 0.6

and p∗K4 = 4.35. Dashed line represents model fit for adsorption in 0.01 M

NaCl ([Cu2+

]t = 1.75 × 10−4

M) with a combination of p∗

K1 = 1.16and p∗K4 = 3.1. However, for 0.1 g/l of adsorbent moderately good fit-

ting (dotted) can be obtained with p∗K1 = 0.70. Dot–dashed line represents

model fit for adsorption in 0.01 M KNO3 ([Cu2+]t = 1.72×10−4 M) with

p∗K1 = 1.2 and p∗K4 = 3.0. Double dot–dashed line represents model fit

for adsorption in 0.5 M NaCl using p∗KintML(1)

=−0.5.

Fig. 10. Variation of fractional adsorption of Zn2+ from different

electrolyte solutions as function of pH. Solid line represents model

fit of adsorption in 0.5 M NaCl ([Zn2+]t = 1.57 × 10−4 M) using

p∗K1 = 0.4. Dashed line represents model fit for adsorption in 0.01 M NaCl

([Zn2+]t = 1.44 × 10−4 M) using p∗K1 = 1.38. Dot–dashed line repre-

sents model fit for adsorption in 0.01 M KNO3 ([Zn2+]t = 1.43×10−4 M)using p∗K4 = 6.1.

Reaction (4) suggests increased adsorption with in-

creased pH, but tends to decrease with increase in concentra-

tion of ligand (e.g., Cl−). Consequently, the adsorption edge

tends to be smeared out and shifts to the higher pH side. Ad-

sorption according to reactions (5) and (6) does not envisage

variation with pH, as no H+ is released, but an increase in ei-

ther L− concentration or M2+ ion concentration (or increase

in adsorbate/adsorbent ratio) will also shift the edge to the

alkaline side. However, if metal–ligand complexation takes

place in the solution phase, ionic strength variation will have




little effect on adsorption [16]. In the case of reaction (6)

increased ligand concentration will drastically reduce the ad-

sorption, as its concentration term will be in the numerator

of the equilibrium constant.

Adsorption according to reaction (7) is favored at lower

pH values where the ligand tends to adsorb more strongly

than the metal ion. Additional adsorption sites are createdwhere the ligand acts as a bridge between the surface and

the cation (type II). Formation of such a complex tends to

increase the metal ion adsorption at lower pH, but decreases

with increase in pH and therefore adsorption edge tendsto be

broad. This is what observed for the adsorption edge of Zn2+

in 0.5 M NaCl solution. The stability sequence of metal–

chlorocomplex (MCl+) is as follows [17]: Ni2+ > Co2+ >

Cu2+ > Zn2+. Therefore, Ni2+ and Co2+ form stronger

chlorocomplexes in 0.5 M NaCl than in 0.01 M NaCl. Given

that nitrate ion does not form any complex in solution, lower

adsorption of Ni2+ and Co2+ in 0.5 M NaCl is due to the

formation of such complex in solution which exhibits loweradsorption than free metal ions [16]. However, in 0.01 M

NaCl solution free metal ion outcompetes solution complex

due to its relatively lower stability, resulting in high ad-

sorption. This is further reflected in the adsorption of Zn2+,

where the Cl− forms ligand-like complexes with the surface

due to lower stability of the ZnCl+ complex in NaCl solu-

tion. The adsorption of Cu2+ is somewhat different as unlike

other metal ions, it forms various hydrolytic species in the

pH range 3–6 that can enhance its adsorption greatly [6].

Therefore, not much difference in adsorption can be noticed

in all the three electrolyte solutions.

Two experiments have also been carried out for the ad-

sorption of Ni2+ and Cu2+ in 0.01 M NaCl solution by

doubling the adsorbate/adsorbent ratio, i.e., using half the

amount of δ-MnO2. It can be seen that whereas in case of

Ni2+ the adsorption edge is more flattened and shifts to

higher pH side, in case of Cu2+ the adsorption edge although

shifts to alkaline side is reasonably steep. This suggests that

in case of Cu2+ there is little or no competition between free

Cu2+ and complexated species for the surface sites, while

in case of Ni2+ metal–chlorocomplex outcompetes Ni2+ for

the surface sites. However, precise prediction of adsorption

in such a complex situation is still very difficult because of

other effects such as manganese ion release, surface hetero-

geneity, etc.

3.3. Determination of surface binding constants of cations

The proton stoichiometry as determined from adsorption

isotherm experiments at different pH values, though not very

precise, are no doubt low in order to suggest that adsorption

does not take place entirely through the formation of either

monodentate-complex-like SOM+ or hydrolytic-complex-

like SOMOH. Initially, an attempt was made to estimate the

intrinsic constant by plotting the logarithm of apparent con-

stant for the formation of the above surface species against

the fraction of surface covered by adsorbed cation ( Γ ad/N s )

and extrapolating to zero adsorption, where the electrosta-

tic factor is considered to be negligibly small [14]. However,

the intrinsic constants obtained in this manner were not help-

ful to obtain a good fit of the experimental data for both

free and complexated metal ions. It was, therefore, consid-

ered that a binding constant such as the one derived from the

Kurbatov-type plot [18], which represents the average of allthe processes involved, would be appropriate in this adsorp-

tion system.

The general equation for the interaction of the metal ion

with an oxide substrate may be expressed by the following

equation:

(8)Mn++ x(SOH)Ke

⇄ [M(SOH)x](n−x)++ xH+.

Neglecting the activity coefficient not only for the metal ion

but also its surface complex the equilibrium constants may

be expressed as follows:

(9)Ke =[M(SOH)x](n−x)+[H+]x

[Mn+][SOH]x .

Taking logarithm and rearranging one obtains the following

relationship:

(10)log [Mn+]ad

[Mn+]soln= log Ke + x

pH+ log[SOH]

.

A plot of left hand side vs {pH + log[SOH]} should yield

linear relationship, the slope and intercept of which give the

values of x and Ke, respectively. We have taken free SOH

into consideration in the present calculation as the amount

adsorbed is not very small compared to the total surface

sites. The results obtained for different electrolyte solutions

from the linear relationships are shown in Table 4. Which

reveal some interesting features. The first thing that may

be noted is that, except Co2+, the proton coefficient (x) in-

creases as the concentration of NaCl solution is decreased

from 0.5 to 0.01 M. In the case of Ni2+, Cu2+, and Zn2+

adsorption, this value is further increased in 0.01 M KNO3,

which is more indifferent than NaCl solution.

On the other hand, Ke decreases with decreased concen-

tration of NaCl and also in 0.01 M KNO3 solution. For the

adsorption of Cu2+, Ke value remains unchanged in 0.01 M

NaCl and 0.01 M KNO3. It may be mentioned here that

when the mode of adsorption changes from more of specificchemical interaction to more of coulombic interaction (in di-

lute electrolyte solution), binding constant tends to decrease.

This does not necessarily imply that adsorption is lower in

the later case and indeed, adsorption in most cases is en-

hanced.

3.4. Modeling of adsorption with respect to pH

Unlike hydrous iron oxide very little work has been done

on the modeling of adsorption of cations, especially heavy

metal ions, onto hydrous MnO2 surface [19–21]. The cited

papers do not deal with the effect of anions, either added in




Table 4

Parameters derived from Kurbatov-type plots of the variation of fractional adsorption of cations with pH (batch titration)

Electrolyte Co2+ (1.81× 10−4 M) Ni2+ (1.81 × 10−4 M) Cu2+ (1.59× 10−4 M) Zn2+ (1.48× 10−4 M)

solution pH (x) − log Ke pH (x) − log Ke pH (x) − log Ke pH (x) − log Ke

range range range range

0.5 M 2.65–4.25 0.38 0.53 2.30–7.12 0.32 0.76a 2.60–6.20 0.58 1.20 2.17–5.62 0.31 0.38

NaCl (0.6)a (0.9)a (0.4)a

0.01 M NaCl 2.10–5.75 0.51 0.60 2.37–5.9 0.54 0.96 2.04–3.48 0.62 1.18 2.80–7.60 0.49 1.40

(0.7)a (1.0)a (1.16)a (1.38)a

0.01 M KNO3 2.05–5.50 0.45 0.65a 2.07–5.85 0.67 1.7 2.03–4.85 0.61 1.15 3.16–4.70 0.70 2.5a

(1.6)a (1.2)a

Note. All batch titration experiments in the presence of Co2+ , Cu2+ , and Zn2+ were carried out using 0.2 g/l of adsorbent. In the case of Ni2+ the amount of

adsorbent used was 0.4 g/l.a Values used for model fitting as p∗Kint

1 .

the system or present as electrolyte medium. Balistrieri and

Murray [22] have ruled out the effect of chloride ion on the

adsorption of Co2+, Ni2+, Cu2+, and Zn2+ onto goethite in

major ion sea water. If no such effect is observed for goethite

with pHpzc lying between 7.2 and 7.5, it is rather more un-likely that electrolyte anion will exert any effect in the case

of δ -MnO2, whose pHpzc lies at 1.5–1.8. Consequently the

same authors [23] have found that 85% of the surface of δ -

MnO2 in sea water is occupied by protons and the remaining

15% by Ca2+, Mg2+, K2+, and Na2+ and not by any chloro-

complex.

The triple-layer model has been used by the earlier re-

searchers [19,21] in their attempts to model adsorption of

cations on δ-MnO2. However, it has been demonstrated in

our previous paper [9] that, the basic Stern model is better

applicable to the δ -MnO2 /electrolyte solution interface than

to the triple-layer model. Therefore, adsorption is consid-ered to take place on the surface plane only. This is evident

from the decreasing value of the apparent constant (pKappM )

of reaction between surface and M2+ so that the difference

between intrinsic and apparent constants (p∗K intM −p∗K

appM ),

which is equal to the electrostatic factor in the following

equations tends to be negative with increase in pH value.The

two most widely used surface complexation reactions are as

follows:

(a) Direct reaction with the surface plane as

monodentate complex

(11)SOH0 +M2+

∗K int

1⇄ SOM++H+(aq),

(12)∗Kint1 = K

app1 exp(F Ψ 0/RT ).

where

(13)Kapp1 =

[SOM+][H+]

[SOH0][M2+].

(b) Adsorption of monohydroxy species onto surface

leading in the neutralization of surface charge

(14)SOH0 +M2++ H2O⇄ SOMOH+ 2H+(aq).

The first proton is released from the coordination layer of

the aquo–metal ion complex ([M(H2O)6]2+

). The combined

equilibrium reactions may therefore be written as

(15)∗K int3 =

[SOMOH][H+]2

[SOH0][M2+].

The species within brackets represent their correspondingactivities in the solution. Adsorption of metal ions under the

above two conditions may be calculated from the following

equations by neglecting the power terms of adsorption den-

sity, as this quantity tends to be small:

[M2+]ad = N s [M2+]tot

N s + [M2+]tot

(16)+[H+]

∗K int1

exp(eΨ 0/ k T )

,

(17)

[M2+

]ad = N s [M2+

]tot

N s +N s[M2+

]tot +

[H+]2

∗K int3

.

The formation of surface complexes involving anion may

take place as follows:

(c) Adsorption of metal–anion complex leading to the

neutralization of surface charge as in the case of (b)

SOH0 +M2++ L−⇄ SOML0 +H+(aq),

(18)∗K intML(1) =

[SOML0][H+]

[SOH0][M2+][L−].

The corresponding adsorption of cation is expressed asfollows:

(19)[M2+]ad =N s[M

2+]T

N s + [M2+]T + [H+]

∗K intML(1)

[L−]

.

The ∗K intML(1) values involving the chlorocomplex are ob-

tained analytically by plotting KappML(1)

against adsorption

densities of Co2+, Ni2+, Cu2+, and Zn2+ in 0.5 M NaCl so-

lution and extrapolating to zero adsorption density are 0.25,

−0.10, −0.5, and −1.8, respectively. But the adsorption val-

ues calculated from Eq. (19) using these intrinsic constants

values are low for Ni2+

up to pH 3.5, whereas for Cu2+

and




Co2+ the values are very low for the entire pH range. In the

case of Zn2+, the calculated values are negligibly small.

An attempt has also been made to model adsorption of

the four metal ions in 0.5 M NaCl according to the following

reactions:

(d) Surface reaction of metal–anion complex without releasing H + in solution

SOH0 +M2++L−⇄ SOHML+,

(20)∗KintML(2) =

[SOHML+]

[SOH0][M2+][L−]exp(F Ψ 0/RT ).

(e) Adsorption of metal–anion complex already formed in

the solution phase

SOH0 +ML+⇄ SOHML+,

(21)∗KintML(3) =

[SOHML+]

[SOH0

][ML+

]

exp(F Ψ 0/RT ).

Both the models yield very low values of adsorption even

at pH above 5.0. Even the combination of models (c) and (d)

or (c) and (e) does not make good fit of the data. Therefore,

anion complexation is not considered as a significant factor

for the decrease in adsorption with increase in ionic strength

of the electrolyte medium. In view of this observation, ad-

sorption of heavy metal ions has been discussed below on

the basis of model reactions (a) and (b).

While adsorption of Ni2+ from 0.5 M NaCl may be mod-

eled as a monodentatesurface complex (SONi+)withp∗K int1

value of 0.76 without the necessity of adding Mn2+ release,

adsorption from 0.01 M NaCl and 0.01 M KNO3 can bemodeled satisfactorily with the combination of SONi+ and

SONiOH surface complexes. However, no satisfactory mod-

eling is possible for the adsorption from 0.01 M NaCl when

0.2 g/l of adsorbent is used. This is possibly because avail-

able sites are restricted compared to the competing solution

species, which is contrary to the basic requirement of the

surface complexation model.

In the case of adsorption of Zn2+ from NaCl solution,

moderate to satisfactory fitting of data can be achieved for

the formation of SOZn+ surface complex together with

Mn2+ release in solution. However, in 0.01 M KNO3 so-

lution fractional adsorption values calculated on the basis of

SOZnOH formation are lower than those of experimentallyfound values even after taking Mn2+ release into considera-

tion.

In case of Cu2+ the adsorption data in all the three elec-

trolyte solutions are best fitted using the combined forma-

tion of SOCu+ and SOCuOH surface complexes along with

Mn2+ release in solution. At higher adsorbate/adsorbent ra-

tio, i.e., using 0.1 g/l δ-MnO2, adsorption data can be mod-

eled with the formation of SO−Cu+ surface complex. But

at pH above 4.0 the calculated values are lower than the ex-

perimental values. Modeling of the adsorption of Co2+ on

δ-MnO2 in all the three electrolyte solutions has been less

satisfactory compared to other metal ions. While in 0.5 M

NaCl and 0.01 M KNO3, solution adsorption takes place as

SOCo+, in 0.01 M NaCl solution, combination of SOCo+

and SOCoOH appears to be the most appropriate for model

fitting.

Hydrous MnO2 is not an ideal substrate for model fit-

ting of adsorption data, because of highly heterogeneous

nature of surface arising from different oxidation statesof Mn, a low-order of crystallinity, and occurrence of ad-

sorbed/occluded alkali metal ion on the surface during its

preparation. Previous attempts [19,21] based on the triple-

layer model are not better than the present work.

4. Conclusions

From the present work on the adsorption of trace metals

on δ-MnO2 the following conclusions can be drawn.

1. Adsorption isotherms of Co2+

, Ni2+

, Cu2+

, and Zn2+

onto δ -MnO2 surface from 0.5 M NaCl solution in the

presence of buffer solution show that adsorption in-

creases with increase in pH. For Cu2+ and Zn2+ adsorp-

tion is very high at pH 6.0. The adsorption isotherms of

all the four cations are non-Langmuirian even at very

low concentrations.

2. The proton stoichiometry derived from the log–log plot

of adsorption vs equilibrium concentration using the

method of Perona and Leckie varies between 0.3 and

0.8 depending upon the cations and their concentration

range up to which linear behavior is followed. Simi-

lar low-proton stoichiometry values have also been ob-

tained from Kurbatov-type plots of adsorption edge.

This suggests that some adsorption takes place without

releasing proton into the solution.

3. The Langmuir-type plots of adsorption isotherm show

breaks in the linear behavior after a certain concen-

tration. The binding constants derived from the initial

linear region increase with decrease in pH and when

extrapolated to pHpzc of the oxide sample give values

closer to those obtained from Kurbatov-type plots.

4. The adsorption of Co2+, Ni2+ and to some extent Cu2+

with respect to pH is generally lowest in 0.5 M NaCl and

highest, in 0.01 M NaCl, while the adsorption in 0.01 M

KNO3 lies between them. In the case of Zn 2+ the ad-sorption in 0.5 M NaCl is highest, at least up to pH 4.0;

in 0.01 M KNO3 it is the lowest, while in 0.01 M NaCl

the adsorption is intermediate between them. It is pos-

tulated that in the case of adsorption of Co2+ and Ni2+

in 0.5 M NaCl, relatively stable chlorocomplexes exhibit

weak metal-like adsorption, whereas in the case of Zn2+

poor stability of chlorocomplex in solution leads to the

adsorption of Cl−, which acts as a bridge between sur-

face and metal ion (ligand-like), especially at lower pH.

5. The adsorption of cations can be modeled either by the

formation of SOM+ type complex or as a combina-

tion of SOM+

and SOMOH type complexes, depending




upon the cation and the nature of electrolyte medium.

The binding constants obtained from the intercepts of

the Kurbatov-type plots are used in such model calcula-

tion.

6. Attempts made to model adsorption on the basis of

surface reaction with metal–chlorocomplexes (MCl+)

have not been successful, as the calculated values of ad-sorption are too low. It is concluded that adsorption of

cations in such a manner occurs only to a very limited

extent compared to free metal ion.

Acknowledgments

The authors are thankful to Dr. Vibhuti N. Mishra, Di-

rector, Regional Research Laboratory, Bhubaneswar, for his

kind permission to publish the paper. One of the authors

(S.S.T.) is grateful to the CSIR, New Delhi, for the award

of Senior Research Fellowship.

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