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8/11/2019 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.
doi:10.1016/j.jcis.2003.07.002
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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
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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.
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14 S.B. Kanungo et al. / Journal of Colloid and Interface Science 269 (2004) 11–21
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 – – – –
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S.B. Kanungo et al. / Journal of Colloid and Interface Science 269 (2004) 11–21 15
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 ,
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16 S.B. Kanungo et al. / Journal of Colloid and Interface Science 269 (2004) 11–21
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.
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S.B. Kanungo et al. / Journal of Colloid and Interface Science 269 (2004) 11–21 17
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
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18 S.B. Kanungo et al. / Journal of Colloid and Interface Science 269 (2004) 11–21
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
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S.B. Kanungo et al. / Journal of Colloid and Interface Science 269 (2004) 11–21 19
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
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20 S.B. Kanungo et al. / Journal of Colloid and Interface Science 269 (2004) 11–21
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
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S.B. Kanungo et al. / Journal of Colloid and Interface Science 269 (2004) 11–21 21
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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