Colloids and Surfaces
A: Physicochemical and Engineering Aspects 207 (2002) 13 – 24
www.elsevier.com/locate/colsurfa
Surface complexation modeling of cadmium adsorption on
gibbsite
R. Weerasooriya a,*, H.K.D.K. Wijesekara a, A. Bandara b
a
Chemical Modeling Project, Institute of Fundamental Studies, Kandy 2000, Sri Lanka
b
Department of Chemistry, Uni6ersity of Peradeniya, Peradeniya, Sri Lanka
Received 8 February 2001; accepted 18 December 2001
Abstract
Cadmium adsorption on gibbsite was examined as a function of pH, background electrolyte concentration, and
adsorbate loading. The adsorption data was quantified by charge distribution multi-site ion complexation (CD
MUSIC) model using reaction stoichiometry given; 2( \ AlOH − 1/2) +Cd2 + = (\ AlOH)x2 Cdy; log K= 6.727. The
charge distribution factor, f is treated as an adjustable parameter in model fitting yielding x =1.2 valence units (vu)
and y= 0.8 vu. © 2002 Elsevier Science B.V. All rights reserved.
Keywords: Cadmium; Gibbsite; Charge distribution multi-site ion complexation model; Langmuir model
1. Introduction
Reactions at the particulate-water interface are
one of the key processes in determining transport
and fate of trace elements in natural environment
[1]. When particulate materials are enriched with
natural organic moieties such as humic substances
(HS), the fate of metal ions is determined to a
large extent by chemical complexation [2]. However, there are situations in which inorganic particulates such as metal hydrous oxides are equally
important or even dominate the overall sorption
process [3]. This is the case for environments in
which very little organic matter is present, as is
* Corresponding author. Tel.: +94-8-232002; fax: + 94-8232131
E-mail address: [email protected] (R. Weerasooriya).
encountered for example in many aquifers or in
clay liners or landfills [4,5]. Unfortunately, mechanistic understanding of metal ions binding on
natural soils, sediments or aquifers is an arduous
task, partly due to inherent complexity of these
solids [6]. Over the past three decades or so, much
attention was, therefore paid to quantify metal
ions binding on well-characterized mineral phases
that dominate in nature. The rationale here is to
consider these discrete mineral phases as ‘building
blocks’ of natural solids in that they could be
combined in such a way to identify dominant
adsorptive surfaces. From this viewpoint substantial chemical data on trace element binding on
iron hydrous oxides has become available (Ref. [7]
and references therein). However, similar data on
gibbsite (and other Al, Si, or Mn) hydrous oxides
that are also common in nature is comparably low
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R. Weerasooriya et al. / Colloids and Surfaces A: Physicochem. Eng. Aspects 207 (2002) 13–24
[8 –16]. Further, none of these studies address
the issue of cadmium adsorption mechanism on
gibbsite.
The aim of this work was therefore, to determine the behavior of cadmium adsorption on
gibbsite over a range of experimental conditions
that are important environmentally. This experimental system was selected primarily due to following factors. (a) Cadmium is a nonessential
element; it is released into the environment
largely from various anthropogenic activities. It
poses serious threat to plants, animals and humans because of their toxicity and persistence.
Hence, the US EPA has categorized it, as a
priority pollutant designating maximum contaminant level goal is 0.005 mg l − 1 [17]. (b) Cadmium exhibits low tendency for polymerization;
therefore it is particularly suited to probe nature
of active surface sites. (c) Cadmium typically
forms weaker aqueous and surface complexes
than do with other heavy metals; therefore Cd
should exhibit high mobility particularly under
slightly acid conditions. (d) Moreover, gibbsite
is a useful analog of naturally occurring aluminum oxides, hydroxides and aluminous silicate
minerals; it offers reactive surface sites when in
contact with surface and ground waters. (e)
Gibbsite is a simple solid that has \Al2OH0
and \AlOH − 1/2 surface sites; only \ AlOH − 1/
2 is found to be reactive when the pH is ranged
4– 10 [12]. Therefore, it provides a convenient
way for a model testing. (f) Recently, the chemical data on gibbsite-solution interface over a
range of experimental conditions has become
available [13–16]; hence, it provides a convenient way to model the interfacial chemistry of
gibbsite– cadmium system. In this work, a series
of experiments were conducted to assess the
cadmium adsorption as a function of pH, adsorbate loading, and electrolytic concentration. Particular attention was paid to determine the
degree of site heterogeneity of cadmium adsorptive centers by non-calorimetric method. After a
critical assessment of existing mechanistic models in light of current requirements, we have selected a charge distribution multi-site surface
complexation (CD MUSIC) model to quantify
cadmium adsorption data [18]. This model emphasizes the importance of the surface structure
and of the charge distribution at the interfaces.
Further, the CD MUSIC model is capable to
account for the finite size of adsorbing ions, and
the site heterogeneity based on crystal faces. The
charge of the adsorbing ions is specifically distributed between two electrostatic planes according to the modified Pauling’s bond valence
theory. Such an approach is not possible in
widely used 2-pK triple-layer model (TLM) or
Frumkin isotherm based models [19–23]. However, in the CD MUSIC modeling the surface is
treated planer. Moreover, the lateral interactions
of surface sites have not been explicitly accounted for.
Further, the CD MUSIC model formulation
requires explicit selection of an electrostatic
model. Presently, we have selected the 1-pK
three-plane model (TPM) since by making several assumptions it can readily be degenerated
into constant capacitance (CCM), Stern layer
(SLM) or diffuse layer models (DLM). However, the 1-pK TPM suggested by Venema et al.
[24] differs from the classical 2-pK TLM with
respect to following factors; outer-layer capacitance, C2 (TLM C2 = 0.2 F m − 2; TPM C2 = 5
F m − 2), and ion pairs placement (TLM, ionpairs in SLM, TPM, in diffused layer). The following additional factors merit the selection of
CD MUSIC model for our calculations. When
compared to 2-pK modeling, the 1-pK approach
significantly reduced the total number of adjustable parameters (e.g. 2-pK TLM, 7 parameters; 1-pK TPM 5 parameters). The intrinsic
acidity constants derived by 2-pK approach are
dependent on electrostatic model used; however
such a limitation is not present in the 1-pK
modeling technique; the acidity constant is simply derived experimentally as pHzpc of the solid.
Further assumptions based on pKi (i= 1 or 2)
are no longer needed (2-pK models are often
constrained by a ZpKi ). The temperature (T)
dependence of surface charging can simply be
predicted by assessing pHzpc. Finally, placing
them in two different planes distinctly separates
the outer- and inner-sphere complexes [18,24].
R. Weerasooriya et al. / Colloids and Surfaces A: Physicochem. Eng. Aspects 207 (2002) 13–24
2. Materials and methods
2.1. Materials
Gibbsite was obtained from ALCOA (Australia). All reagents were prepared in de-ionized,
distilled water that was de-gassed by boiling and
then purging with N2 (purity 99.996%). All pH
adjustments were made either with 0.541 M
HNO3 or 0.600 M NaOH. Particular attention
was paid to store NaOH under N2 to prevent
possible CO2 contamination. All experiments were
performed in a glove box that was purged with
99.996% N2 to minimize atmospheric contamination of CO2 according to following flushing cycle;
thrice prior to use; twice after sample change, and
twice daily to minimize any slowly diffused CO2
into the glove box. The gas samples inside the
glove box was checked after each flushing cycle
with the GC/TCD detector, and was always below 5 pM. The 1000 mg l − 1 (8.897 mM) cadmium
stock solution (AAS grade) used was obtained
from Mercks, Germany. Table 1 shows the
parameters (e.g. specific surface area, site density,
intrinsic acidity and electrolyte binding constants,
formation constants of cadmium species used in
calculations.
2.2. Methods
mined as a function of pH, ionic strength, and
adsorbate loading. Unless otherwise mentioned,
the system temperature was always kept at 298 K.
A 500 ml of 2 g l − 1 gibbsite suspensions was
prepared in 0.001 M NaNO3, and let it equilibrated for 2 h. Always the suspension pH ranged
between 7 and 8. Before the addition of calculated
aliquot of 0.890 mM Cd2 + to reach desired metal
ion concentration in the range of 0.77–7.74 mM,
the system pH was first adjusted to 3.5 with
0.541 M HNO3. Ten milliler aliquots were then
sampled into 40 ml capacity polystyrene cen-
Table 1
Surface parameters and reaction stoichiometries used in CD
MUSIC modeling of cation adsorption
Parameter/reaction
Value
Source (Reference)
Parameter
Surface area (m2 g−1)
Site density (sites nm−2)
13
8.5
[15]
[15]
Stern layer capacitance
C1 (F m−2)
C2 (F m−2)
2.58
5
[12]
Protolysis reaction
\AlOH+1/2
“
2
8.70
[15]
\ AlOH
−1/2
2.2.1. Adsorption kinetics
Chemical kinetics of cadmium adsorption on
gibbsite was determined in 0.01 M NaNO3 at pH
5.5. The pH adjustment was made with HNO3.
Relevant experimental details are well documented [16], hence only an outline is given.
Briefly, a 250 ml of batch solution was prepared
with 2 g l − 1 gibbsite, and it was spiked with stock
cadmium solution to reach a final concentration
at 0.78 mM. At pre-defined time intervals, samples
(5 ml aliquots) were withdrawn into a syringe,
which was fixed, to a 0.45 m disposable filter
cartridges for immediate solid-solution separation. The filtrate was acidified at pH B1 for
cadmium analysis.
2.2.2. Adsorption edges
Cadmium adsorption on gibbsite was deter-
15
+H+
s
Ion pair formation
−
\AlOH−1/2
+H+
s
s +NO3 “ 7.975
[15]
\AlOH+1/2
−NO−
2
3
\AlOH−1/2
+Na+
s
s “
[15]
−0.396
\ AlOH−1/2−Na+
Aqueous phase reactions
Cd2++2NO3− =Cd(NO3)20 0.00
Cd2++2OH− =Cd(OH)20
7.70
Cd2++3OH− =Cd(OH)3− 8.99
Cd2++4OH− =Cd(OH)4−2 8.71
Cd2++5OH− =Cd(OH)5−3 8.07
Cd2++6OH− =Cd(OH)6−4 7.19
2Cd2++2OH− =Cd2(OH)+2 7.60
4Cd2++4OH− =Cd4(OH)4 28.08
Cd2++NO3− =CdNO3+
0.31
Cd2++OH− =CdOH+
3.90
H++OH− =H2O
−1.40
Na++OH− =NaOH0
−0.20
[25]
[25]
[25]
[25]
[25]
[25]
[25]
[25]
[25]
[25]
[25]
[25]
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R. Weerasooriya et al. / Colloids and Surfaces A: Physicochem. Eng. Aspects 207 (2002) 13–24
trifuge tubes for successive addition of variable
volumes of 0.600 M NaOH to adjust the pH
between 3 and 10 at 0.5 pH unit intervals. Once
the upper pH value was reached (normally 10),
addition of 5 M NaNO3 was made to adjust
background ionic strength to 0.01 M. Soon after
the system pH was decreased to 3.5 with 0.541
M HNO3 and samples were taken again after
adjusting the pH between 3.5– 10. The whole process was repeated to obtain sample aliquots at 0.1
M NaNO3. Samples were equilibrated in a temperature regulated water bath orbital shaker
(Eyela, B 603, Japan). The pH of the samples was
measured after an equilibration period of 24 h.
After phase separation by centrifugation (14 000
rpm) at regulated temperature for 5– 10 min., and
followed by membrane filtration (0.45 m), the
filtrates were acidified for metal analysis. The
degree of reversibility of cadmium adsorption was
also examined. The methodology followed remained same except following changes. The ionic
strength was kept at 0.01 M. The spiking of 0.77
mM Cd was made for a well-equilibrated gibbsite
suspension at pH 8. The pH adjustments were
done with 0.541 M HNO3 before sampling 10 ml
of solid suspension into centrifuge tubes for 24 h
equilibration.
2.2.3. Adsorption isotherms
Cadmium adsorption isotherms were constructed in 0.01 M NaNO3 as a function of pH.
Experimental protocols remained essentially same
as stated in the section under adsorption edges
except following changes. The 2 g l − 1 gibbsite
suspension was equilibrated for 2 h at a desired
pH. Afterwards, a 10 ml portion of samples was
transferred into centrifuge tubes. The addition of
different quantities of cadmium into these tubes
was made after matching the solution pH.
The total concentration of cadmium was varied
between 0.1 and 1000 mM. At the end of
24 h equilibration period, the pH of the stirred
suspension was measured, and regulated at desired value typically for 0.5–1 h till a stable pH
reading is achieved. Subsequently solid– solution
phases were separated to determine solute concentration.
2.2.4. Analyses
All pH measurements were made using Ross
combination pH electrode and auto chemistry
analyzer (Orion 960) in a well-stirred solid suspension. Cadmium concentration was measured either with flame (Cd \2 mM) or graphite furnace
(GF) (Cd B 2 mM) atomic absorption spectroscopy (AAS) using an atomic absorption spectrometer (AA) equipped with a GF and an auto
sampler (AA GBC 933AA, GF GBC GF 3000,
and GBC PAL 3300). Both flame and GF measurements were made at 228.8 nm. In the GF-AA
pyrolitic graphite coated tubes were used with L’
van platforms. The atomization program consisted of following steps: drying (two steps): step1 :
temperature 90 °C ramp time 10 s, hold time 10 s;
step2 : (120 °C, 20 s, 10 s); pyrolysis (400 °C, 1 s,
5 s); and atomization (1800 °C, 0.8 s, 0.8 s) with
the inert purge gas flow arrested and flushing
(2000 °C, 1 s, 5 s) (in atomization step, the inert
purge gas flow was arrested. Air flow rate was 4.5
ml min − 1 in other steps).
2.2.5. Data modeling
The surface of gibbsite consists of several structurally different functional groups. These groups
differ due to the surface oxygens on gibbsite being
singly and doubly coordinated to aluminum. The
structure of gibbsite forms stacked sheets of
linked octahedrons of Al(OH)3. The octahedrons
are composed of Al3 + coordinated to six OH ions
yielding + 0.5 (vu) according to Pauling’s theory.
Two aluminum ions are needed to neutralize one
OH− in the gibbsite structure. Thus in the crystal
structure all hydroxyls are doubly coordinated.
The planar faces of gibbsite have only Al2OH0
groups whereas the edge faces have either Al2OH0
or AlOH − 1/2 groups. As calculated by Hiemestra
et al. [12] the protonation of \Al2OH occurs
when pH 0. Similarly the de-protonation of \
Al2OH sites occur when pH 12. In the pH range
that we considered the \ Al2OH are inert. Hence,
only \ AlOH − 1/2 sites were treated as active. All
calculations were performed using ECOSAT-FIT
utility [25,26]. Very recent release of ECOSAT has
incorporated FIT optimization algorithm to extract binding and other constants from the adsorption measurements.
R. Weerasooriya et al. / Colloids and Surfaces A: Physicochem. Eng. Aspects 207 (2002) 13–24
17
3. Results and discussion
3.1. Adsorption kinetics
Interactions of solids with water involve a combination of chemical processes; i.e. surface hydration and hydroxylation, adsorption and
desorption, diffusion. It is known that surface
hydration is a slow process [7,21], which gradually
adds adsorption sites to the system that result an
enhanced ion adsorption over prolong periods of
equilibration. Although it is difficult to infer information about a given process from a simple
kinetic experiment such data is used to determine
the time required for first appearance of ion equilibration. Therefore, the adsorption kinetics was
studied in order to determine the time required
for equilibrium to be achieved. It is evident from
the data in Fig. 1 (A) that the cadmium uptake by
gibbsite was sluggish reaching an apparent
plateau within first 6– 7 h. Previously, a two-step
process has been used to characterize the retention of metal ions on hydrous oxides; a rapid
adsorption preferably due to the external surfaces
followed by a slow retention due to intra-particle
diffusion along the micro pore walls utilizing internal surface area [27]. In the slow adsorption
process, with time the amount of metal ions
sorbed to the oxide gradually increased due to
inter-particle surface diffusion and surface hydration over extended equilibrium periods. In our
experiments, within the first 24 h equilibration
period, such an observation was not observed for
cadmium adsorption on gibbsite. Within this
equilibration period, as shown in Fig. 1 (A), the
time-dependent sorption curve is essentially flat.
This indicates that the surface sites present in
external surfaces of gibbsite govern the cadmium
adsorption within first 24 h equilibration period.
3.2. Hysteresis
The hysteresis between cadmium adsorption
and desorption curves was also examined. As
shown in the Fig. 1 (B), the cadmium adsorption
on gibbsite was completely reversible with pH
within 12 h; in other words, within the limits of
experimental error, decreasing the pH yielded a
Fig. 1. (A) Variation of GCd (cadmium adsorption density) as
a function of equilibration time. Initial conditions used were;
[Cd2 + ]= 7.78 mM; pH 7; T= 298 K; 0.01 M NaNO3.
Similar trends were observed both at system pH 8 or 9
(data not given). (B) Variation of GCd as a function of pH. In
desorption, gibbsite – cadmium complex was first formed at pH
8.0 after equilibration for 24 h. Subsequently, the metal – gibbsite complex was treated with 0.01 M NaNO3 at different pH
for 24 h. Relevant data is shown by opened symbols. Filled
symbols represent adsorption data obtained by varying pH
from 4 to 10. Measurements were made in triplicate to determine error bars. Experimental conditions employed were;
gibbsite = 2 g l − 1 and T= 298 K.
18
R. Weerasooriya et al. / Colloids and Surfaces A: Physicochem. Eng. Aspects 207 (2002) 13–24
desorption curve that follows the adsorption
curve. Based on the data given in Fig. 1 (A and
B), therefore, it was concluded that cadmium
adsorption on gibbsite was dominated by a sluggish, reversible-surface mechanism when 24 h
equilibration period was employed.
3.3. Adsorption isotherms
Cadmium adsorption density, GCd, was determined for different pH in 0.01 M NaNO3. The
results obtained are shown as adsorption
isotherms in Fig. 2 (A). At all instances, the GCd
increased with equilibrium cadmium concentration approximating to a slope one. If cadmium
were polymerized, then the slope would be steeper
than one. Slopes greater than one were not observed in any of the isotherms investigated [23].
This implied that only monomeric cadmium species reacts with the surface.
Solution chemistry of cadmium suggests that it
can form several chemical species. Relative abundance of these species was calculated using thermodynamic data (Table 1). In the calculation, the
precipitation of cadmium was not taken into account due to following factors. The solubility
product, Ksp of Cd(OH)2 is 10 − 14.35 at 298 K [28].
Cd(OH)2 precipitation occurs when [Cd2 + ]\ 1
mM. In our system, the total [Cd2 + ]B 1 mM
which shows an unsaturated condition with respect to Cd(OH)2 precipitation. Fig. 2 (B) shows
the distribution of various cadmium species as a
function of pH. It is evident from the diagram,
that only free cadmium ion is abundant in solution when pH ranges 4–9. Both CdOH+ and
Cd(OH)20 become significant when pH \ 10.
Fig. 2. (A) Variation of Gads with equilibrium [Cd2 + ] in 0.01 M NaNO3 at 298 K as a function of pH. Symbols represent
experimental data. Error analyses were done on triplicate experimental data. Solid/broken lines show CD MUSIC/TPM model
calculations carried out by using parameters given in Tables 1 and 2. The slope, m was calculated by linear regression of log[Gads]
vs. log[Cd] curves. When m : 1 adsorption data can be interpreted in terms of Langmuir equation. Otherwise (m 01), adsorption
is governed by Freundlich equation. B: Variation of cadmium species as a function of pH in 0.01 M NaNO3. Initial [Cd] =8.7 mM.
The system was unsaturated with respect to cadmium. Hence, no solid precipitates were assumed.
R. Weerasooriya et al. / Colloids and Surfaces A: Physicochem. Eng. Aspects 207 (2002) 13–24
Table 2
Langmuir isotherm constants for cadmium adsorption at pH
7.5, 7.1, and 6.0
pH
K 9 error
(m3 mol−1)
Gmax 9 error (mol m2−)
Rr 2
7.5
7.1
6.0
140 9 64
30.0 9 15
21.8 913
4.2E−049 8.6E−05
4.9E−049 8.5E−05
7.3E−069 4.2E−07
0.97
0.98
1.00
In all cases, the temperature was fixed at 298 K. Isotherm data
were calculated from a non-linear least square fit method and
the related correlation coefficients (R 2) were also given in the
table. R 2 values were almost 0.98 or higher which indicate a
very good mathematical fit. However, the error estimates are
high, hence no attempt was made to report any data below pH
6.
However, our working pH window was ranged
from 4 to 9.5. Therefore, it can be concluded that
under the experimental conditions imposed in this
study, only free cadmium ion is present in a
significant proportion for surface interactions.
Further, as shown in Fig. 2 (A) the general
shape of the isotherms can be modeled with a
Langmuir equation:
Gads =
k[M]Gmax
(1+ k[M])
(1)
where Gads and [M] are the concentrations of
adsorbate at the surface (mol m − 2) and in solution (M), respectively, k is the equilibrium constant for the adsorption– desorption process, and
Gmax is the amount of adsorbate required to form
a mono-layer. The Langmuir model assumes that
the energy of surface sites is equal (e.g. homogeneous surface).
The parameters k and Gmaxwere estimated
simultaneously, using a non-linear least square fit
of the experimental data, and the optimized values are given in Table 2. Within the limits of error
estimates, both k and Gmax have shown an apparent increase with the pH. The k represents an
empirical affinity constant that depends on Gads.
The increase in Gmax with the pH is attributed to
an increase in the fraction of surface sites that
favor cadmium adsorption. The pHzpc of gibbsite
is 8.7. When pHBpHzpc, the majority of surface
sites are charged positively. As the pHzpc is approached, a significant proportion of these sites
19
will be neutralized and a small number will acquire a negative charge before the pHzpc is
reached. Given the adsorbing species (e.g. Cd2 + )
is positively charged, adsorption will occur most
readily on neutral or negative surface sites. The
increase in Gmax with pH is therefore due to an
increase in the fraction of neutral sites available
for adsorption, as pHzpc [29]. However, dominance of surface sites of single class (\ AlOH − 1/
2) does not provide conclusive evidence for
homogeneous surface sites. Because of surface
roughness, on an atomic scale there will be a
difference in reactivity among sites with the same
coordination [30]. These effects when combined
should generate a surface with a wide range of
adsorption energies that can subsequently result
in a heterogeneous surface. Accordingly, we have
assumed that at least for the experimental conditions imposed in the study, that the \ AlOH − 1/2
are homogeneous.
3.4. Adsorption edges
Fig. 3(A, B and C) shows the variation of
cadmium adsorption density (GCd) as a function
of ionic strength and pH at different initial cadmium concentrations. The GCd has increased
(from 1% to almost 100%) with the pH increase
from 4 to 9.5. Within the experimental error, no
difference in cadmium adsorption is observed for
values below pH 6. Above this pH value, the
adsorption edge occurs within a narrow pH range
(less than 2 pH units) when [Cd]initial 0 8 mM.
When initial [Cd]=77.4 mM, the broadening of
adsorption edge occurs between pH 6.2 and 8.5.
In all cases, it illustrates a typical sigmoid curve
characteristic of d-group metals; similar trends
were observed for lead and copper adsorption on
gibbsite [15,16].
The inflection point of the cadmium adsorption
edges, which correspond to the pH of 50% adsorption (pH50%) as determined by d2G/dpH2 was
always found to be within 7.50–8.00 for a 100fold variation of [Cd]initial. The pH of the 50%
adsorption (i.e. pH50%) was 7.40 for Cd 0.77 mM,
7.70 for Cd mM and 7.80 for Cd 77 mM. As shown
in the data, the position of the adsorption edge
shifts slightly to a higher pH as the initial metal
20
R. Weerasooriya et al. / Colloids and Surfaces A: Physicochem. Eng. Aspects 207 (2002) 13–24
concentration increases. At a given experimental
condition, when the initial cadmium concentration is set so that the adsorption density data falls
Fig. 3.
in the linear region of the isotherms (shown in
Fig. 2 (A)), the position of the adsorption edge (as
determined by pH50% values) showed virtually no
shift with the adsorbate loading. When the relative position of the edge shifts with the adsorbate
loading, it implies that the initial experimental
conditions are set to fall in the non-linear region
of the isotherm. In all cases, whether the initial
cadmium concentration falls in the linear or nonlinear region, a maximum adsorption density,
Gcdmax, was reached when pH \7.5. Further, the
Gcdmax value was directly proportional to the initial [Cd]initial. Such proportionality cannot, however, be deduced when Gcd B Gcdmax. Therefore, it
is concluded that cadmium adsorption on gibbsite
over the entire pH range examined exhibits a
non-proportional behavior for the variation of
adsorbate loading and pH.
Relative shifts of the adsorption edges with
respect to a given reference point for variation of
ionic strength have often been used to infer bonding mechanism(s) of surface complexes [31]. The
adsorption density of weakly bonded complexes
has showed a marked effect for the variation of
ionic strength, yielding a distinctly separated adsorption edges. Such an observation cannot however be made when ions are strongly bonded.
Further, demarcation of a boundary between
weak vs. strong surface complexes is somewhat
arbitrary yielding ambiguous results for ions that
exhibit ‘intermediate’ affinity for surfaces. Furthermore, it is important to note that the relative
shifts of adsorption edges always do not yield
Fig. 3. Variation of Gads as a function of pH, ionic strength
and adsorbate loading. In all cases 2 g l − 1 gibbsite was used.
(A) Initial [Cd2 + ]= 0.77 mM; (B) 7.74 mM; (C) 77.4 mM.
Replicate measurements were made only for data set (B). All
measurements were made at 298 K. All lines (both solid and
broken) show calculated data by using CD MUSIC model.
Only cadmium – gibbsite bidentate complex was considered.
Charge distribution of Cd2 + between 0- and 1-planes was
calculated according to following equations; Zz0 =nHzH +
fzMe and Dz1 =(1 −f )zMe + Smizi. The nH is the change of
protons on the surface ligands involved in the surface reactions; zH is valence of protons zMe charge of central ion in
surface complex; mj is number of ligands positioned in 1-plane;
zi is the charge of the ligands in 1-plane; f is the charge
distribution factor calculated according to the modified bond
valence theory.
R. Weerasooriya et al. / Colloids and Surfaces A: Physicochem. Eng. Aspects 207 (2002) 13–24
conclusive evidence into the binding mechanism.
For example, species with large log K will have a
smaller effect on ionic strength than with a low
log K, irrespective of the way of bonding. In the
present case, as shown in Fig. 3, the pH50%, (the
pH corresponds to 50% adsorption) differs only
by 0.4 pH units for a 100-fold variation of
NaNO3 concentration. This points to a strong
interaction with the surface, possibly yielding an
inner-sphere complex losing the sheath of H2O
molecules present in the hydration shell. Molecular scale information provided by in situ spectroscopic methods is needed to precisely define
bonding mechanism and surface speciation of adsorbed ions at a mineral–water interface. Such
data for our cadmium– gibbsite system is in progress. In the meantime, a-Al2O3 was selected as an
analogous solid for gibbsite due to following
points. Both solids contain AlO8 octahedra; the
oxygens are closest-packed hexagonally and two
thirds of interstitial sites occupied by Al atoms.
In-situ extended X-ray absorption fine structure
spectroscopy (EXAFS) measurements have shown
that cadmium formed inner-sphere complex at the
a-Al2O3 –water interface [32]. The macroscopic
adsorption data presented here for gibbsite– cadmium system also supports this notion.
3.5. Proton exchange ratio, rH/Cd
This is defined as a measure of surface protons
released as a result of cadmium adsorption. The
exact mechanism of surface protons release is,
however a controversy to date. It is assumed that
21
a part of the rH/Cd is attributed to direct release of
surface proton(s) upon the metal ion adsorption
that subsequently increased the surface charge.
Release of additional protons through coulombic
repulsion compensates any changes of surface
charge [7]. Recently, Rietra et al. [33] have shown
that the electrostatic interactions are dominant
factor in governing this ratio and not the stoichiometry of the reaction. Table 3 shows the rCd/H
data that was determined with a pH-stat approach. When [Cd]initial/[Cd]absorbed 1 and pH \
5, the rCd/H always converges to a non-integer
value between 1.33 and 1.5 (Table 3), which
shows the release of, more than a proton per
metal ion adsorbed.
3.6. Data modeling
The cadmium adsorption data presented in Fig.
2 (A) and Fig. 3 was modeled using CD MUSIC
model. The specific reasons for the selection of
this model were discussed under the introduction
section. The CD MUSIC model requires seven
parameters to be determined. As listed in Table 1,
following data of gibbsite-solution interface is
available in the literature; C2, outer-layer capacitance was assigned a value of 5 F m − 2. The
inner-layer capacitance C1 is an adjustable
parameter, which was optimized to give best description of surface charge data. As outlined under the Materials and Methods section, gibbsite
possess only AlOH − 1/2 sites that are active. Site
homogeneity is assumed, at least for the stringent
conditions imposed in the study, as a result of
Table 3
Release of protons upon cadmium adsorption on gibbsite at different fixed pH values
pH
7.10
7.40
8.00
a
GCd, mmol m−2
0.32
0.63
1.58
[H+]released/[Cd+2]adsorbed
Experimental 9S.D.a
Nb
R 2c
Calculatedd
1.3290.09
1.51 9 0.01
1.5590.30
3
4
3
0.90
0.89
0.99
1.40
1.60
1.65
Standard deviation.
Number of samples used.
c
Linear correlation coefficient.
d
All calculations were done using CD MUSIC model.
b
22
R. Weerasooriya et al. / Colloids and Surfaces A: Physicochem. Eng. Aspects 207 (2002) 13–24
Table 4
Optimized surface complexation constants of Cd(II) determined by pCd titration data
NaNO3, M
Log KCda
b
R2
c
0.001
0.01
0.1
7.118
6.476
6.326
0.99
0.91
0.99
6.569E−09
3.331E−08
7.369E−09
Weighted av. Log KCd
6.727
RMSE
All experiments were conducted in 0.01 M NaNO3 at 298 K.
Model calculations were carried out using CD MUSIC model.
Weighted averages were calculated using log K= Swi (log K)I
where wi = (1/| log K)I and s log K is the standard deviation
calculated from ECOSAT-FIT for i h data set.[7,25,26].
a
Surface complex [(\AlOH)2x Cdy]2+ = 2(\AlOH−1/2)+
Cd2+ log KCd.
b
R 2 equals to one in a perfect fit.
c
RMSE: residual mean of sum of squares.
satisfactory fits by Langmuir model. Further,
ionic strength dependence data of cadmium adsorption suggest an inner-sphere bonding mechanism. In theory, cadmium can form both
monodentate and bidentate surface complexes
with active \ AlOH − 1/2. Bargar et al. [34] em-
ployed an empirical model based on the modified
bond valence concept to predict the relative stability of different kinds of adsorbed species on the
aluminum surface in an aqueous solution. According to this model, the charge saturation of the
central atom (presently the oxygen) is calculated
(Sl0…m ). A given surface complex is thermodynamically stable when Sln…m B z, where z is the
charge of the central atom. If the sum is greater
than 2.0 vu, desorption of ion from the surface
should occur. Presently, the charge of the central
ion (e.g. oxygen) in the bidentate cadmium complex is unsaturated leading to a stable configuration. In the case of mondentate complex, the
charge of the surface oxygen is over-saturated,
which means that this complex is thermodynamically unstable. Hence the reaction stoichiometry
corresponding to bidentate cadmium complex was
chosen for CD MUSIC modeling. In order to
account for the finite size of adsorbing ion, the
CD MUSIC model requires the charge of the
surface complex is distributed in the interface
between o- and b-planes. Presently we have considered the charge distribution factor as a fitting
parameter. As shown in Table 4, the log K value
Fig. 4. CD MUSIC model calculations of different surface species of cadmium and \ AlOH − 1/2 in 0.01 M NaNO3. In the
calculation the following initial conditions were imposed; solid content = 2 g l − 1. (A) Initial [Cd2 + ]= 0.77 mM (B) [Cd]=7.74 mM;
(C) [Cd]=77.4 mM NaNO3 = 0.01M. All other parameters used were given in Table 1 and Table 2.
R. Weerasooriya et al. / Colloids and Surfaces A: Physicochem. Eng. Aspects 207 (2002) 13–24
of bidentate cadmium surface complex is optimized using the cadmium binding data given in
Fig. 3 (B). Intrinsic binding constants of cadmium
surface complexes were obtained from the adsorption edge data shown in Fig. 3 (B). These model
parameters were subsequently utilized to quantify
adsorption edges constructed at initial Cd2 + 0.77
and 7.78 mM, and adsorption isotherms shown in
Fig. 2 (A). As shown in the Fig. 2(A) and Fig. 3
by solid and broken lines, the model calculations
agree reasonably well with experimental data.
Finally, different surface species of gibbsite and
cadmium were calculated using CD MUSIC
model, and the results obtained are given in Fig.
4. When compared to [(\ AlOH)21.2Cd0.8], free
[\ AlOH − 1/2] and [\AlOH1/2] represent a major
portion of the surface sites. According to crystallographic data, the maximum site density of gibbsite was estimated as 14 mmol m − 2. In all
experiments, maximum Gcdnever exceeded 3.0
mmol m − 2. Therefore, gibbsite surface is unsaturated with respect to the adsorbate loading when
initial [Cd2 + ]B7.78 mM.
4. Conclusions
Surface sites of gibbsite can be treated homogeneous for cadmium ion binding. Only cadmium
bidentate surface complex, \Al(OH)21.2Cd0.8 was
used in CD MUSIC modeling. When [Cd]initial/
[Cd]absorbed 1 and pH \ 5, the rH/Cd was ranged
between 1.2 and 1.6. This is attributed to a release
of more than a proton per metal ions adsorbed.
Acknowledgements
We thanked International Foundation for Science (IFS, Sweden) for offering travel supports
under research grant no. AF 2407-2 and gifting
Orion Auto-chemistry Analyzer to our group
(Grant No. A.F.2407). Meindert (WAU, The
Netherlands) is thanked for supplying ECOSATFIT (ver. 4.7) to our group. The comments made
by anonymous reviewer have enhanced
manuscript quality.
23
References
[1] W. Stumm, Chemistry of the Solid – Water Interface, Wiley, New York, 1992.
[2] J. Buffle, Complexation Reactions in Aquatic Systems:
An Analytical Approach, Ellis Horwood, London, 1988.
[3] J.A. Davis, J.A. Coston, D.B. Kent, C.C. Fuller, Environ.
Sci. Technol. 32 (1998) 2820.
[4] K. Spitz, J. Moreno, A Practical Guide to Groundwater
and Solute Transport Modeling, Wiley, New York, 1996.
[5] B.K. Schroth, G. Sposito, Environ. Sci. Technol. 32
(1998) 1404.
[6] A.W.P. Vermeer, J.K. Mcculloch, W.H. van Riemsdjik,
L.K. Koopal, Environ. Sci. Technol. 33 (1999) 3892.
[7] D.A. Dzombak, F.M.M. Morel, Surface Complexation
Modeling: Hydrous Ferric Oxide, Wiley, New York,
1990.
[8] F.J. Hingston, A.M. Posner, J.P. Quirk, J. Soil. Sci. 23
(1972) 176.
[9] M.B. McBride, Soil. Sci. Soc. Amer. J. 49 (1985) 843.
[10] M.B. McBride, Clays Clay Min. 38 (1982) 21.
[11] T. Hiemstra, W.H. van Riemsdijk, M.G.M. Bruggenmert,
Neth. J. Agric. Sci. 36 (1987) 281.
[12] T. Hiemstra, W.H. van Riemsdijk, Coll. Surf. 59 (Section
A) (1991) 7.
[13] W.N. Rowlands, W. O’Brien, R.J. Hunter, V.J. Patrick, J.
Coll. Interf. Sci. 188 (1997) 325.
[14] T. Hiemstra, H. Young, W.H. van Riemsdijk, Langmuir
15 (1999) 5942.
[15] R. Weerasooriya, B. Dharmasena, D. Aluthpatabendi,
Coll. Surf. 170 (Section A) (2000) 65.
[16] R. Weerasooriya, D. Aluthpatabendi, H.J. Tobschall,
Colloid Surf., Section A, 189 (2001) 131.
[17] US Environmental Protection Agency, Drinking Water
Standards, 1995. Available from http://www.epa.gov.
[18] T. Hiemstra, P. Venema, W.H. van Riemsdjik, J. Coll.
Interf. Sci. 181 (1996) 45 – 59.
[19] L.K. Koopal, W.H. van Riemsdjik, M. Groffery, J. Coll.
Interf. Sci. 116 (1987) 197.
[20] J.A. Davis, D. Kent, in: M.F. Hochela, A.F. White
(Eds.), Mineral – Water Interface Geochemistry, American
Geophysical Union, Washington DC, 1990.
[21] J. Westall, in: W. Stumm (Ed.), Aquatic Surface Chemistry, Wiley, 1987.
[22] E. Gileade, E. Kirowa-Eisner, J. Penciner, Interfacial
Electrohemistry: An Experimental Approach, AddisonWesley, MA, 1972.
[23] H. Tamura, N. Katayama, R. Furuchi, J. Coll. Interf. Sci.
195 (1997) 241.
[24] P. Venema, T. Hiemstra, W.H. van Riemsdjik, J. Coll.
Interf. Sci. 183 (1996) 515.
[25] M. Keizer, W.H. van Riemsdijk, ECOSAT version 4.7,
Equilibrium Calculations of Reactions and Transport,
Department of Soil Science, Wagenigen Agricultural University, The Netherlands, 1999.
[26] D. Kinniburgh, FIT User Guide Technical Report WD/
93/23, British Geological Survey, Nottingham, 1993.
24
R. Weerasooriya et al. / Colloids and Surfaces A: Physicochem. Eng. Aspects 207 (2002) 13–24
[27] P. Trivedi, L. Axe, Environ. Sci. Technol. 34 (2000)
2215.
[28] A.E. Martell, R.M. Smith, Critically Selected Stability
Constants of Metal Complexes, National Institute of
Standard Technology, USA, 1998 database version 5.
[29] B.B. Johnson, Environ. Sci. Technol. 24 (1990) 112.
[30] K.C. Hass, W.F. Schneider, A. Curioni, W. Andreoni,
Science 282 (1998) 265.
[31] K.F. Hayes, J.O. Leckie, J. Coll. Interf. Sci. 115 (1987)
564.
[32] C. Papelis, G.E. Brown Jr., G.A. Parks, J.O. Leckie,
Langmuir 11 (1995) 2041.
[33] R.P.J.J. Rietra, T. Hiemstra, W.H. van Riemsdjik,
Geochim. Cosmochim. Acta 63 (1999) 3009.
[34] J.R. Bargar, S.N. Towle, G.E. Brown Jr., G.A. Parks, J.
Coll. Interf. Sci. 185 (1997) 473.