See discussions, stats, and author profiles for this publication at: https://www.researchgate.net/publication/284353583
Kinetics of the Reduction of Colloidal MnO2 by Citric Acid in the Absence and
Presence of Ionic and Non-ionic Surfactants
Article in BioInorganic Reaction Mechanisms · January 2005
DOI: 10.1515/IRM.2005.5.3.151
CITATIONS
READS
4
357
3 authors, including:
Syed Mohammad Shakeel Iqubal
Zaheer Khan
Ibn Sina National College for Medical Studies
Jamia Millia Islamia
55 PUBLICATIONS 113 CITATIONS
234 PUBLICATIONS 3,507 CITATIONS
SEE PROFILE
SEE PROFILE
Some of the authors of this publication are also working on these related projects:
Antimicrobial Susceptibility testing View project
Synthesis characterization and biological evaluation on the metal complexes of novel N-substituted phthalimide ligand View project
All content following this page was uploaded by Syed Mohammad Shakeel Iqubal on 29 January 2019.
The user has requested enhancement of the downloaded file.
Inorganic Reaction Mechanisims, Vol. 5, pp. 151-166
Reprints available directly from the publisher
Photocopying permitted by license only
© 2005 Old City Publishing, Inc.
Published by license under the OCP Science imprint,
a member of the Old City Publishing Group.
Kinetics of the Reduction of Colloidal MnO2
by Citric Acid in the Absence and Presence
of Ionic and Non-ionic Surfactants
KABIR-UD-DIN1, *, S.M. SHAKEEL IQUBAL1 and ZAHEER KHAN2
1
Department of Chemistry, Aligarh Muslim University, Aligarh 202 002, India
2
Department of Chemistry, Jamia Millia Islamia, Jamia Nagar, New Delhi 110 025, India
Received on Nov 11, 2003; Revised manuscript received Sept 11, 2004; Date accepted Sept 20, 2004.
The kinetic results of the reduction of water soluble colloidal manganese dioxide by citric
acid in the absence and presence of ionic and non-ionic surfactants are presented. Irrespective
of the medium, the reaction is first- and fractional-order, respectively, in [oxidant] and
[reductant]. The reduction is accelerated by increase in the concentrations of added Mn(II)
and non-ionic Triton X-100 but anionic SDS has no effect. The Arrhenius and Eyring
equations were valid for the reaction over a range of temperatures and different activation
parameters have been evaluated. A plausible mechanism of the reaction is proposed.
Keywords: Colloidal MnO2; citric acid; kinetics; reduction; Triton X-100; SDS
INTRODUCTION
involvement as active autocatalysts in many
permanganate oxidations [10-12].
Surfactants are referred to as amphiphilic,
amphipathic, heteropolar or polar/nonpolar
compounds due to the characteristic of possessing
distinct hydrophobic (water-repelling) and
hydrophilic (water-loving) regions in their molecules.
The interest in using surfactants as reaction media is
that they affect rates, products and, in some cases,
stereochemistry of the reactions [13-15]. Studies of
chemical reactions in micellar media could provide
understanding even about the reactions taking place at
Perez-Benito and his coworkers [1] found first
time quantitatively that water soluble colloidal
manganese dioxide, prepared from reduction of
aqueous potassium permanganate by sodium
thiosulphate under neutral condition [2,3], was
perfectly transparent and stable for several months.
Manganese dioxide (as aqueous suspension) has been
used as an oxidizing [4-7] and catalytic [8,9] agent of
inorganic/organic compounds. The transparent sols of
manganese dioxide too are of importance due to their
*Corresponding author. E-mail: kabir7@rediffmail.com
151
152
KABIR-UD-DIN et al.
interfaces [14-18]. Electron-transfer processes in
micellar systems can be considered as models to get
insight into electron transport occurring in biological
phenomenon [19].
The reduction of colloidal manganese dioxide by
organic acids has received some attention [2,3,20-23]
but that of citric acid has not yet been studied. As the
redox chemistry of citric acid plays an important role in
human nutrition under neutral conditions, it was thought
that break down of biologically important citric acid
using this oxidant (colloidal MnO2) would be worth
studying. The main objective of the present
investigation is to elucidate the mechanism of the redox
reaction between citric acid and colloidal MnO2 both in
the absence and presence of surfactants.
EXPERIMENTAL
Materials
Potassium permanganate, sodium thiosulphate and
citric acid were E. Merck (India) reagent grade
chemicals. The solutions were prepared in deionized
and doubly distilled water. KMnO4 solution was boiled
and filtered into a dark glass bottle for storage and was
standardized by titration against oxalate, whereas
sodium thiosulphate solution was standardized by
titration against potassium dichromate. Water soluble
colloidal MnO 2 was prepared using the method of
Perez-Benito et al. [1].
Kinetic Measurements
The kinetics of MnO2-citric acid redox reaction was
carried out in a three-necked reaction vessel fitted with a
double-walled spiral condenser (to prevent
evaporation). The reaction vessel was immersed in a
water bath thermostated at 30±0.1 °C unless stated
otherwise. The reaction volume was always 50 cm3 and
citric acid was used in excess with respect to colloidal
MnO2. The progress of the reaction was monitored
spectrophotometrically for ~80% completion. The
absorbance for disappearance of colloidal MnO2 was
measured at 390 nm with Baush & Lomb Spectronic20 spectrophotometer. The pseudo-first-order rate
constants (kobs or kψ, s-1) were estimated from the slopes
of the conventional ln(absorbance) versus time plots.
Other details of the kinetic procedure were the same as
described previously [24-27].
RESULTS AND DISCUSSION
The water-soluble colloidal MnO2 obtained from the
reaction represented by Eq. (1) was characterized by
means of the UV-vis spectra of standard KMnO4 and its
reaction product with sodium thiosulphate. The
spectrum of KMnO4 solution
8MnO4- + 3S2O32- + 2H+ →
8 MnO2 + 6SO42- + H2O
(1)
possesses an absorption band located at λmax = 530 nm
(Fig. 1). This spectrum changed by the addition of
sodium thiosulphate, when the band at 530 nm
gradually disappeared with the appearance of a single
broad band of high intensity at 390 nm (Fig. 1). From
these results and previous observations [1,28] it is
confirmed that the new stabilized spectrum is that of
water soluble colloidal MnO2.
(A) Reaction in Absence of Surfactants
The log(A390) versus time plot (Fig. 2, inset) shows
that the reaction of citric acid and colloidal MnO2
proceeds in two stages and that the first stage
(noncatalytic) is relatively slower than the second
(autocatalytic). The kinetic curves (Fig. 2) show
inflection and the kobs values (for both the stages) were
estimated from such curves. It is observed that extent
of noncatalytic path depends upon the reaction
conditions (see Fig. 2(a-d) for the effect of substrate
concentration; similar effects were observed with
increase in [MnO2] as well as temperature). Kinetics
of the reaction was, therefore, investigated at several
initial reactant concentrations (as the increase in
temperature resulted in disappearance of the
noncatalytic path, the kinetic runs were performed at
30 °C — at this temperature the progress of the
reaction was neither fast nor slow) and temperatures
both in the absence and presence of surfactants and
the results are summarized in Tables I and II.
REDUCTION OF COLLOIDAL MnO2 BY CITRIC ACID
153
FIGURE 1
Absorption spectra of KMnO4 (1.6 x 10-4 mol dm-3, ●) and of the reaction product of KMnO4 (8.0 x 10-5 mol dm-3 ) and Na2S2O3 (3.0 x 10-5
mol dm-3, ).
FIGURE 2
Plots of log(absorbance) versus time for the reduction of colloidal MnO2 by citric acid. Reaction conditions: [MnO2] = 8 x 10-5 mol dm-3,
[citric acid] = 16(a), 24(b), 40(c), 56 x 10-4 mol dm-3(d), temperature = 30 °C. Inset – plot of log(absorbance) versus time for curve (a).
154
KABIR-UD-DIN et al.
REDUCTION OF COLLOIDAL MnO2 BY CITRIC ACID
155
TABLE I
Dependence of rate constants on the factors influencing the reduction of colloidal MnO2 by citric acid.
To asses the effect and to find out the order with
respect to [MnO 2 ], the k obs1 and k obs2 values
(pertaining, respectively, to noncatalytic and
autocatalytic pathways) were determined at different
initial MnO2 concentrations in the range 2.0 - 10.0 x
10-5 mol dm-3 at constant [citric acid] (=16 x 10-4 mol
dm-3) and temperature (= 30 °C). It was observed that
as the initial [MnO2] increases the values of both kobs1
and kobs2 decrease (Table I). It is well established that
pseudo-first-order rate constants are independent of
the initial reactant concentration. The abnormal
behaviour probably is due to possible flocculation of
the colloidal particles. Gum arabic is known to
stabilize colloidal MnO2 particles in solution [3,29].
Therefore, some experiments were also performed in
presence of this water-soluble polysaccharide. The
autocatalytic pathway was not observed in presence
of gum arabic (Table I). Also, when gum arabic was
added to the reaction mixtures retardation in kobs1 was
observed (Table I). This type of ambiguity was also
observed in many permanganate reactions
[3,12,30,31].
To investigate the effect of acid concentration on
rate, the reaction was studied as a function of [HClO4]
(0.0 to 30.0 x 10–4 mol dm–3) at fixed [MnO2] (8 x 10–5
mol dm–3) and [citric acid] (16 x 10–4 mol dm–3) at 30
°C. It was observed that the reaction rate increased
markedly with increasing [HClO4] (Table I). The plots
of kobs1 and kobs2 versus [HClO4] were straight lines
(Fig. 3) with positive intercepts on the y-axis,
indicating acid- independent and acid- dependent
reaction paths being involved in the reduction of
colloidal MnO 2 by citric acid. Further, doublelogarithmic plot of k obs1 and [HClO 4 ] yielded a
straight line with slope of 0.19 ( r = 0.9604 )
indicating the order with respect to [HClO4] to be
fractional for the noncatalytic pathway (see inset of
Fig.3 where kobs1 versus [H+]0.19 is shown to yield a
good straight line).
At fixed [MnO 2 ] (= 8 x 10 -5 mol dm -3 ), the
dependence of kobs1 and kobs2 on [citric acid] were also
determined at different reductant concentrations. The
results are depicted graphically in Fig. 4. The values
of rate constants (kobs1 and kobs2) are given in Table II.
The plots of log kobs versus log [citric acid] are linear
(Fig. 4(B)) with positive slopes (0.49, r = 0.9790,
kobs1; 0.28, r= 0.9685, kobs2). Activation parameters
(E a, ∆H ≠ and ∆S ≠) are believed to provide useful
information regarding environment in which chemical
reactions take place. Therefore, a series of kinetic
runs were carried out within the temperature range 20
to 40 °C at fixed [citric acid] (= 16 x 10-4 mol dm-3) and
[MnO2] (= 8 x 10-5 mol dm-3 ). The kobs (Table I) were
found to fit the Arrhenius and Eyring Eqs. (2) and (3):
kobs = A exp(-Ea/RT)
kobs = (kBT/h) exp(-∆H≠/RT) exp(∆S≠/R)
(2)
(3)
where the symbols have their usual meanings. The nonlinear least squares method was used to obtain the
values of parameters which are recorded in Table I.
156
KABIR-UD-DIN et al.
Mechanisms satisfying the above requirements are given in Schemes 1 and 2:
(i) for the noncatalytic pathway:
(4)
(5)
(6)
(7)
(8)
(9)
(10)
SCHEME 1
(ii) for the autocatalytic pathway:
(11)
(12)
(13)
(14)
SCHEME 2
REDUCTION OF COLLOIDAL MnO2 BY CITRIC ACID
157
FIGURE 3
Plots showing the effect of HClO4 on the pseudo-first order rate constants. Reaction conditions: [colloidal MnO2] = 8 x 10-5 mol dm-3, [citric
acid] = 16 x 10-4 mol dm-3, temperature = 30 °C. Inset – plots showing the effect of hydrogen ion concentration on the pseudo-first order rate
constant. Reaction conditions: [colloidal MnO2] = 8 x 10-5 mol dm-3, [citric acid] = 16 x 10-4 mol dm-3, temperature = 30 °C.
In Scheme 1, Eq.(4) represents adsorption of the
active species of citric acid on the surface of the colloidal
MnO2. After adsorption, the species(C1) undergoes
electron transfer process leading
to the formation of 2.
oxo- glutaric acid, radical (COOH) and Mn(II) in a rate
determining step. In the subsequent fast. step, as depicted
in Eq. (6), MnO2 further reacts with C OOH and gives
Mn(III) and CO2 as reaction products. From Scheme 1
mechanism, the following rate law may be derived:
(15)
where k'I = k1Kc1, k''II = k2Kc2Kc1.
And, for the first-order rate constant:
(16)
In Scheme 2, the first step is the complex
formation between the adsorbed species and Mn(II).
The redox decomposition of this complex in the
subsequent rate-determining. step gives rise to 2-oxoglutaric acid, Mn(III) and C OOH radical (Eq. (12)).
The free radical will give CO2 in accordance with Eq.
(6). Manganese(III) is a strong oxidant and is unstable
with respect to disproportionation. However, we tried
to monitor formation of Mn(III) at 470 nm [32] but
failed to detect any build up of the species during the
course of the reaction.
In order to confirm the involvement of Mn(II)
(product of noncatalytic reaction pathway, Eq. (7)) in
the autocatalytic path (Eq. (11)), the same
experiments were performed by adding Mn(II). The
rate constants, obtained as a function of [Mn(II)] with
other variables remaining constant, were found to
increase with increasing [Mn(II)] (Table I). The
results are shown graphically in Fig. 5 . The plot of
log kobs1 versus log[Mn(II)] resulted in two straight
portions with slopes =0.08 and 0.61, suggesting that
158
KABIR-UD-DIN et al.
TABLE II
Values of first-order rate constants at various [citric acid]a.
the reaction is zero and fractional-order dependence with
respect to [Mn(II)] at lower and higher concentrations,
respectively. On the other hand, the reaction shows
sigmoid dependence on [Mn(II)] for the autocatalytic
reaction pathway (Fig. 5). In the presence of Mn(II),
Scheme 1 is, therefore, modified as Scheme 3.
In the presence of Mn(II) there is a competition
between the citric acid and Mn(II) to react with the
colloidal MnO 2. Thus, we may conclude that the
values of kobs1 (Table I, see Mn(II) effect above or
Scheme 3) are the sum of the paths I and II. As the
contribution of path I cannot be completely ruled out
[28], the exact dependence of kobs1 on [Mn(II)] cannot
be predicted . From the above reasoning we can
expect that additional reaction (11) would make the
autocatalytic path more complicated and obtaining an
equation for kobs2 would be of no use.
As far as the sigmoid dependence on [Mn(II)] is
SCHEME 3
concerned (Fig. 5) for the autocatalytic reaction
pathway, the concentration of Mn(III) increased first
and then started to decrease due to the Mn(III)-citric
acid reaction(Eq.(7)), whereas that of Mn(II)
increased, passed through maximum, decreased, and
finally showed a slow increase. The final conversion
of Mn(III) to Mn(II) seemed to occur when Mn(IV)
was not present in the system any longer [3,28].
(B) Reaction in Presence of Surfactants
In order to see any type of interaction between
ionic surfactants (CTAB and SDS) and the colloidal
MnO 2 , a series of experiments were designed in
presence of varying concentrations of the surfactants
at constant reactant concentrations. Due to
flocculation, the redox reaction between colloidal
MnO 2 and citric acid could not be followed in
REDUCTION OF COLLOIDAL MnO2 BY CITRIC ACID
159
TABLE III
Dependence of rate constants on [TX-100] for the reduction of colloidal MnO2 by citric acida.
presence of cationic CTAB which possesses a positive
charge (colloidal MnO2 is known to be stabilized in
aqueous solution by adsorption of anions [1,11,12,33]
imparting negative charge to MnO2 particles [1]). The
observation of anionic SDS micelles producing no
effect on the reaction rate is understandable in view of
the repulsion between the anionic SDS micelles and
negatively charged sols.
The effect of adding different amounts of non-ionic
surfactant TX-100 to a solution containing constant
[MnO2] (= 8 x 10-5 mol dm-3) was seen at 30 °C. As
[TX-100] increased, the absorbance of colloidal MnO2
decreased and the two are related by Eq. (17):
log(A390) = – 0.0363 log [TX-100] – 0.281 (17)
The effect of varying [TX-100] upon the rate of
colloidal MnO2 reduction by citric acid was studied at
constant [citric acid] (= 16 x 10-4 mol dm-3), [MnO2] (=
8 x 10-5 mol dm-3) and temperature (= 30 °C). The
observed data (Table III) are shown graphically in Fig. 6
as rate constant –[surfactant] profiles.
To see the effects of [oxidant], [reductant] and
temperature, and to confirm whether or not the
aqueous medium mechanism is operative in micellar
TX-100, a series of kinetic runs were carried out at
constant [TX-100] (= 15 x 10 -3 mol dm -3). The k ψ
values, obtained as a function of different variables,
are summarized in Table IV. The behavior upon the
variation of above parameters were found to be
identical, i.e., same order dependence on [oxidant]
and [reductant] as in the absence of TX-100, which
clearly indicates that the same mechanisms (Schemes
1 and 2) are being followed in both the media.
The observed effect of TX-100 on kψ (Table III and
Fig. 6) is catalytic up to certain [TX-100]; thereafter, an
inhibitory effect follows. The catalytic effect may be
explained in terms of the mathematical model proposed
160
KABIR-UD-DIN et al.
TABLE IV
Dependence of rate constants on the factors influencing the reduction of colloidal MnO2 by citric acid in presence of TX-100 (= 15 x 10-3
mol dm-3)
by Tuncay et al. [23], according to which
log kψ1 = 0.2146 log [TX-100] – 2.2831 (r=0.9667) (18)
(20)
log kψ2 = 0.0779 log [TX-100] – 2.1299 (r=0.9833) (19)
A linear relationship was also found by using the
following equations:
(21)
REDUCTION OF COLLOIDAL MnO2 BY CITRIC ACID
161
FIGURE 4
Plots showing the effect of citric acid on the pseudo-first order rate constants (A) and logkobs versus log[citric acid] (B). Reaction conditions:
[colloidal MnO2] = 8 x 10-5 mol dm-3, temperature = 30 °C.
162
KABIR-UD-DIN et al.
FIGURE 5
Plots showing the effect of Mn(II) on the pseudo-first order rate constants. Reaction conditions: [colloidal MnO2] = 8 x 10-5 mol dm-3, [citric
acid] = 16 x 10-4 mol dm-3, temperature = 30 °C.
TABLE V
Comparison of second-order-rate constants (kII) for the reduction of colloidal MnO2 by different reductants at 25 °C.
REDUCTION OF COLLOIDAL MnO2 BY CITRIC ACID
163
FIGURE 6
Plots showing the effect of TX-100 on the pseudo-first-order rate constants. Reaction conditions: [citric acid] = 16 x 10-4 mol dm-3, [colloidal
MnO2] = 8 x 10-5 mol dm-3, temperature = 30 °C.
The fulfilment of the above relations can be seen
in Fig. 7. The values of a and b were calculated
from the slopes and intercepts of Fig. 7(B) (a and b
= 7.62 x 10 2s, 3.4561 mol dm –3s and 4.89 x 10 2s,
0.4902 mol dm–3 s, for kψ1 and kψ2, respectively).
Probable Role of TX-100
Adsorption of non-ionic surfactants and gum
arabic (a protective colloid) on the surface of the
colloidal particles is a well known phenomenon. As a
result, gum arabic stabilizes the colloidal MnO2 [29]
whereas non-ionic surfactants enhance the dispersion
stability [34]. Therefore, both should have the same
influence on the reaction rate. As pointed out earlier,
the effects of the two are opposite to one another, i.e.,
TX-100 had a catalytic and gum arabic an inhibitory
effect. These results indicate that adsorption is not the
only factor responsible to explain the catalytic effect
observed with TX – 100 on the reduction of MnO2 by
citric acid. In addition, the role of other factors, e.g.,
hydrogen bonding, properties of interfacial water
(known to be less polar but more structured than bulk
water), differences in stabilization of the initial and
transition states by surfactant molecules, reaction
rates in the bulk and micellar-pseudo phase, etc.,
cannot be ruled out completely.
As far as the role of TX-100 is concerned, hydrogen
bonding between this nonionic surfactant (I) and the
reactants may play an important role. Citric acid
possesses no hydrophobic character and has 3-COOH
groups. Hydrogen bonding may occur between the
–COOH groups of citric acid and the ether-oxygen of
the polyoxyethylene chains of TX-100. Due to the
164
KABIR-UD-DIN et al.
FIGURE 7
Plots of logkψ versus log[TX-100] (A) and (1/kψ - k obs) versus 1/[TX-100] (B). Reaction conditions: [citric acid] = 16 x 10-4 mol dm-3, [colloidal
MnO2] = 8 x 10-4 mol dm-3, temperature = 30 °C. The data belong to the part up to which the effect of TX-100 was catalytic (cf. Fig. 6).
REDUCTION OF COLLOIDAL MnO2 BY CITRIC ACID
presence of a number of donor groups in one TX100 molecule, multiple H-bonding may take place
and the number of bound citric acid molecules
increase. In this surfactant, the lengths of the
hydrophobic and hydrophilic parts are comparable
and have significant amount of water in the outer
shell. The hydrogen bonding between MnO 2 sols
and hydrophilic part (polar ethylene oxide) of the
TX-100 can not be ruled out either (let us call it
adsorption!). Therefore, the associated MnO 2 and
citric acid with TX-100 (through hydrogen bonding)
seem responsible of facilitating the reaction; this
might be the role of TX-100 towards the observed
catalysis. This surfactant thus helps in bringing the
reactants closer, which may orient in a manner
suitable for the redox reaction followed by
rearrangement of TX-100 molecules.
The decrease in kψ at higher [TX-100] (> 15 x
10 -3 mol dm -3 ) could be due to ‘dilution effect’:
continuous increase in [TX-100] produces
micelles and, progressively more and more
substrate (citric acid) gets associated to the
micellar phase. This segregation deactivates the
substrate since citric acid in one micelle cannot
react with MnO2 (onto which TX-100 is
adsorbed). A point worth noting is that even at
[TX-100] = 50 x 10 -3 mol dm -3, the k ψ’s are still
higher than the values observed in bulk water
(Fig. 6); this proves beyond doubt that TX-100 is
still playing its role in catalyzing the reaction.
Finally, to confirm Tuncay’s propositions [23]
of hydroxyl ions bonded to colloidal MnO 2 being
the active points for substrate adsorption on the
colloidal surface, the reactivity of different
reductants (citric, oxalic, formic and lactic)
towards colloidal MnO 2 has been compared (Table
V). Based on their reactivity, the reductants can be
ordered as: citric > oxalic > formic > lactic, which
clearly indicates that the reaction rate increases
with increase of the number of hydroxyl group in
the substrate molecule. As four hydroxyl groups
are contained in each citric acid molecule,
reaction rate is expected to be higher in
comparison to others (Table V).
165
REFERENCES
[1] Perez-Benito, J. F., Brillas, E. and Pouplana, R. (1989), Inorg.
Chem. 28, 390.
[2] Perez-Benito, J. F. and Arias, C. (1992), J. Colloid Interface Sci.
149, 92.
[3] Perez-Benito, J. F., Arias, C. and Amat, E. (1996), J. Colloid
Interface Sci. 177, 288.
[4] Sharma, T. C., Lal, A. and Saksena, V. (1976), Bull. Chem. Soc.
Jpn. 49, 288.
[5] Basak, B. and Malati, M. A. (1977), J. Inorg. Nucl. Chem. 39,
1081.
[6] Kienzle,F. (1983), Tetrahedron Lett. 24, 2213.
[7] Taniguchi, S. Bull. (1984), Chem. Soc. Jpn. 57, 2683.
[8] Abd El-Salaam, K. M. (1975), Z. Phys. Chem. (Frankfurt am
Main) 95, 139.
[9] Coughlin , R.W. and Matsui, I. (1976), J. Catal. 41, 108
[10] Freeman, F. and Kappos, J. C. (1985), J. Am. Chem. Soc. 107,
6628.
[11] Perez-Benito, J. F. and Lee, D. G. (1985), Can. J. Chem. 63,
1275.
[12] Perez-Benito, J. F. and Arias, C. (1991), Int. J. Chem. Kinet.
23, 717.
[13] Fendler, J.H. and Fendler, E. J. (1975) Catalysis in Micellar
and Macromolecular Systems (Academic Press, New York).
[14] Bunton, C. A. (1979), Catal. Rev. Sci. Eng. 20, 1.
[15] Bunton, C. A. (1997), J. Mol. Liq. 72, 231. 17
[16] Khan, M. N. (1997), Colloids Surf. 127, 211.
[17] Kabir-ud-Din, Salem, J. K. J., Kumar, S. and Khan, Z. (1999)
J. Colloid Interface Sci. 215, 9.
[18] Kabir-ud-Din, Hartani, K. and Khan, Z. (2001), Colloids Surf.
A 193, 1.
[19] Fendler,J.H. (1982) Membrane Mimetic Chemistry (John
Wiley and Sons, New York).
[20] Bradley, J. and Van Praagh, G. (1938), J.Chem. Soc. 1624.
[21] Xyla, A. G., Sulzberger, B., Luther, G. W., Hering, J. G., Van
Cappellen, P. and Stumm, W. (1992), Langmuir 8, 95.
[22] Pimienta, V., Lavabre, D., Levy, G. and Micheau, J. C. (1994),
J. Phys. Chem. 98, 13294.
[23] Tuncay., M., Yuce, N., Arlkan, B. and Gokturk, S. (1999),
Colloids Surf. A 149, 279.
[24] Khan, Z., Rafiquee, M. Z. A. and Kabir-ud-Din. (1997),
Transition Met. Chem. 22, 350.
[25] Khan, Z. and Kabir-ud-Din. (1999), Int. J. Chem. Kinet. 31,
409.
[26] Kabir-ud-Din, Akram, M. and Khan, Z. (2001), Colloids Surf.
A 178, 167.
166
KABIR-UD-DIN et al.
[27] Khan, Z., Raju and Kabir-ud-Din.(2003), Colloids Surf. A 225,
75.
[31] Wiberg, K.B. and Stewart, R. (1955), J. Am. Chem. Soc. 77,
1786
[28] Khan, Z., Raju, Akram, M. and Kabir-ud-Din. (2004), Int. J.
Chem. Kinet. 36, 359.
[33] Jaky, M. and Zrinyi, M. (1993), Polyhedron 12, 1271.
[29] Tompkins, F. C. (1942), Trans. Faraday Soc. 38, 131.
[30] Senent, S. and Cuadrado,S. (1961), An. Quim.Ser. B. 57, 11.
View publication stats
[32] Macartney, D.H. and Sutin, N. (1985), Inorg. Chem. 24, 3403.
[34] Cummins, P. G., Staples, E. and Penfold, J. (1990), J. Phys.
Chem. 94, 3740.