Micellar effect on pentavalent vanadium oxidation of formaldehyde to formic acid in aqueous acid media at room temperature

Article abstract of DOI:10.1007/s11164-014-1635-4

The kinetics of oxidation of formaldehyde by pentavalent vanadium in 3.0 mol dm^-3 H₂SO₄, at 313 K, under pseudo first-order conditions [Formaldehyde]_T [V(V)]_T, have been studied by UV-visible spectrophotometry. Two representative non-functional surfactants (sodium dodecyl sulfate, SDS, and polyoxyethylene octyl phenyl ether, TX-100), at concentrations above and below their critical micelle concentrations (CMC), were used as micro-heterogeneous catalysts in this oxidation. The reaction rate and selectivity strongly depend on the surfactant used, and sometimes on surfactant concentration also. The CMC for both surfactants in aqueous media were determined by spectrofluorimetry, from the sharp change in fluorescence intensity. In contrast with TX-100, SDS was an excellent catalyst of oxidation of formaldehyde by vanadium(V) in aqueous micellar media, leading to the corresponding oxidation product. Formic acid was detected by ¹H NMR spectroscopy. Formation of aggregates by the catalytic surfactants under the reaction conditions was studied by scanning electron microscopy. Dynamic light scattering was used to characterize shape changes during oxidation, by monitoring changes in the hydrodynamic diameter (D_h = 2R_h, where R_h is hydrodynamic radius) of aggregates. A mechanism proposed for this micelle-catalysed slow oxidation reaction was entirely supported by our experimental results.

Full text of DOI:10.1007/s11164-014-1635-4

Res Chem Intermed
DOI 10.1007/s11164-014-1635-4
Micellar effect on pentavalent vanadium
oxidation of formaldehyde to formic acid
in aqueous acid media at room temperature
•
•
•
Pintu Sar Aniruddha Ghosh Rumpa Saha
Bidyut Saha
Received: 2 January 2014 / Accepted: 23 March 2014
Ó Springer Science+Business Media Dordrecht 2014
Abstract The kinetics of oxidation of formaldehyde by pentavalent vanadium in
3.0 mol dm-3 H2SO4, at 313 K, under pseudo first-order conditions [Formalde-
hyde]T ꢀ [V(V)]T, have been studied by UV–visible spectrophotometry. Two
representative non-functional surfactants (sodium dodecyl sulfate, SDS, and poly-
oxyethylene octyl phenyl ether, TX-100), at concentrations above and below their
critical micelle concentrations (CMC), were used as micro-heterogeneous catalysts
in this oxidation. The reaction rate and selectivity strongly depend on the surfactant
used, and sometimes on surfactant concentration also. The CMC for both surfactants
in aqueous media were determined by spectrofluorimetry, from the sharp change in
fluorescence intensity. In contrast with TX-100, SDS was an excellent catalyst of
oxidation of formaldehyde by vanadium(V) in aqueous micellar media, leading to
1
the corresponding oxidation product. Formic acid was detected by H NMR spec-
troscopy. Formation of aggregates by the catalytic surfactants under the reaction
conditions was studied by scanning electron microscopy. Dynamic light scattering
was used to characterize shape changes during oxidation, by monitoring changes in
the hydrodynamic diameter (Dh = 2Rh, where Rh is hydrodynamic radius) of
aggregates. A mechanism proposed for this micelle-catalysed slow oxidation
reaction was entirely supported by our experimental results.
Electronic supplementary material The online version of this article (doi:10.1007/s11164-014-1635-
4) contains supplementary material, which is available to authorized users.
P. Sar ꢁ A. Ghosh ꢁ R. Saha ꢁ B. Saha (&)
Homogeneous Catalysis Laboratory, Department of Chemistry, The University of Burdwan,
Golapbag, Burdwan 713104, WB, India
e-mail: b_saha31@rediffmail.com
R. Saha
K. G. Engineering Institute, Bishnupur, Bankura, WB, India
123

P. Sar et al.
Keywords Catalysis ꢁ CMC ꢁ Micro-heterogeneous ꢁ SDS ꢁ TX-100 ꢁ
Formaldehyde ꢁ Oxidation ꢁ Kinetics
Introduction
Transition metals induce reductive and oxidative transformations, depending on their
oxidation states [1]. Higher-oxidation-state transition metals, for example chro-
mium(VI), manganese(VII), lead(IV), cobalt(III), and vanadium(V) are of the most
effective oxidizing agents for oxidation of organic compounds [2] and use of these
substances as oxidizing agents is of high theoretical and industrial interest. They
behave as either one equivalent or two equivalent oxidants, depending on the reaction
conditions [3]. Among the metals mentioned above, Cr(VI) [4, 5], Pb(IV), and Co(III)
[6] are highly toxic and carcinogenic; in contrast, Mn(VII) and vanadium(V) are
relatively non-toxic and non-carcinogenic. The standard redox potential of V(V)/
V(IV) is 1.00 V [7] at 298 K, information which suggests selection of vanadium(V) as
an oxidant for a variety of organic substrates. Vanadium chemistry is mainly
dominated by the ?4 and ?5 oxidation states, and the ?5 oxidation state is well
known for behaving as an oxidizing agent. We therefore selected vanadium(V) as
oxidant for oxidation of formaldehyde in micellar medium.
Much work has been conducted on the kinetics of oxidation of organic
substrates [2, 8–21] and inorganic compounds [22] by vanadium(V), but no
systematic kinetic study of the oxidation of formaldehyde by vanadium(V) has
been reported. Vanadium(V) acts as an oxidizing agent and catalyses a variety of
catalytic reactions, for example the epoxidation of allylic alcohols, the hydrox-
ylation of alkene, the oxidation of sulfide, and halide oxidation [10, 23]. Many
sugars (sucrose, glucose, fructose, maltose, etc.) and sugar alcohols (mannitol,
sorbitol, etc.) are oxidized by pentavalent vanadium (V2O5) [10]. The unique one-
electron transfer redox process of vanadium enables a variety of novel oxidation
and reductions of organic compounds.
Formic acid is an essential organic compound, principally used as a preservative
and antibacterial agent in livestock feed. It is also used as a pesticide to treat and
control mites that infest honey bee hives. Formic acid is produced in different ways
[24–27]. Oxidation of formaldehyde by a variety of oxidizing agents [25] is one of
the best methods of production of formic acid, and is widely used in industry.
The kinetics of oxidation of formaldehyde by Cr(VI), Co(III) [28], and Ce(IV)
[29] ions have been studied and reported elsewhere. Because we are interested in
kinetic study of the oxidation of formaldehyde, we very recently reported an instant
and facile oxidation reaction of formaldehyde by Ce(IV) [29].
NH4VO3
HCHO
HCOOH
H+ , 313 K
NH4VO3 H+
313 K
,
,
SDS or TX-100
123

Oxidation of formaldehyde by pentavalent vanadium
In this study we focussed mainly on the kinetics of the oxidation of toxic
formaldehyde to industrially important formic acid by vanadium(V) in aqueous and
micellar media.
We observed that vanadium(V) oxidized formaldehyde to formic acid in
3.0 mol dm-3 H2SO4 medium very slowly, and the reaction takes a long time for
completion. Therefore, a suitable catalyst is required to increase the speed of this
oxidation process. However, the requirement for a catalyst of this reaction with high
catalytic activity at room temperature is a challenge. From a survey, it was found
that efficient catalytic oxidation of formaldehyde is uncommon. The difficulty in
oxidizing formaldehyde arises because in aqueous medium it exists in a hydrated
form (methylene glycol). Water is everywhere in the environment, and the hydration
equilibrium for formaldehyde lies far to the right (K * 2.28 9 103) [30].
OH
H
H
Keq
C
H2 O
C
O
OH
H
H
However, despite extensive studies of the kinetics of oxidation of formaldehyde
by vanadium(V), use of surfactants as catalysts in similar studies is not well known.
Surfactants have widespread industrial and technological applications, for example
in pharmaceutical processes, emulsification, wetting ability, and encapsulation and
solubilising of hydrophobic drugs in the human body [31, 32], biodegradability, and
bioaccumulation, etc. Surfactant self-assemblies, for example micelles, vesicles,
etc., in aqueous solution are a major area of interest. The relationship between
surfactant structure and morphology of the self-aggregate, and the different
molecular interactions which determine the properties of the aggregates are now
beginning to be understood [33, 34]. The catalytic effects of surfactants and micelles
on oxidation kinetics have emerged as powerful means of detection of reaction
pathways, and warrant special attention [35].
Surfactants are molecules containing both hydrophilic and hydrophobic parts.
Above a particular concentration, surfactant molecules form different kinds of
aggregate in water, for example micelles, vesicles, and lyotropic liquid crystal
phases [36]. Micelles and vesicles differ not only in their structures but also in their
sizes; whereas the diameter of a micelle is 3–7 nm, multi-lamellar and uni-lamellar
vesicles vary in size between approximately 1,000 nm and 25 nm, respectively [37].
An amphiphilic aggregate in aqueous medium will afford a reaction environment
different from that of bulk water, leading to a kinetic effect on a reaction, and may
be capable of catalysing or inhibiting reactions [38, 39]. The phenomenon of rate
enhancement in a particular reaction by use of surfactants is often called micellar
catalysis. The suitable choice of surfactant can lead to a 5–1,000-fold increase in
reaction rate compared with the same reaction in the absence of surfactant [40]. As
expected, nonionic micelles have only a small effect on the rate of a catalytic
reaction [41] whereas cationic micelles interact with negatively charged substrates
and anionic micelles interact with positively charged substrates. This may be
because of strong electrostatic interactions at the binding sites. Our results support
123

P. Sar et al.
this hypothesis. It has been reported that all surfactant monomers aggregate can
increase the speed of reactions quite effectively (KW = 7.6 9 10-6 s-1) [42].
Homogeneous micellar catalysis by variety of surfactants have been reported [18–
21, 28, 43–45] by several workers but there is no substantial evidence in support of
rate enhancement.
Dynamic light scattering (DLS) measurements are useful for measurement of
particle size distributions, but extracting information about such geometrical features
as shape and stability is difficult. We therefore monitored the stability of our micellar
system under our reaction conditions by use of scanning electron microscopy (SEM).
We used both SEM and DLS measurements to determine the shape, and size
distribution of micelles.
Work on the kinetics of oxidation of formaldehyde by transition-metal ions in
perchloric acid media has been reported elsewhere [28]. Here, we performed this
oxidation in sulfuric acid media both in the presence and absence of surfactant
solution. The purpose of this study was to investigate the kinetics of oxidation of
formaldehyde by vanadium(V) in the presence and absence of different micellar
catalysts, and to investigate mechanistic aspects of micellar effects on the kinetics.
The presence of surfactants was found to accelerate the reaction. Therefore, to
investigate the catalytic effect of vesicles and micelles of the corresponding
surfactant we used two experimental techniques, SEM and DLS, as already stated. It
is of special interest to understand how their morphology and size vary in the
presence of both the substrate, i.e., formaldehyde, and the oxidant used. This work
shows that our procedure is very much simpler, faster, and more efficient than other
catalytic oxidation systems used for this purpose.
Experimental
Materials and methods
Formaldehyde, SDS, Triton X-100, and Congo red (SRL, Mumbai, India), ammonium
metavanadate (Fischer Scientific, Mumbai, India), sulfuric acid, CDCl3, and
acrylonitrile (Merck, Mumbai, India), and other chemical reagents were used without
further purification. All the chemicals used in this work were of AR grade. Double-
distilled water was used to prepare reaction mixtures.
A stock solution of the oxidant was prepared by dissolving accurately weighed
ammonium metavanadate (NH4VO3, white solid) in 3 M sulfuric acid to produce a
yellow solution. The stock solution so prepared was standardized against a standard
ferrous ammonium sulfate solution, with N-phenylanthranilic acid as indicator, as
reported in the literature [17]. A stock saturated solution of formaldehyde was
prepared in double-distilled water. Surfactant (SDS and TX-100) solutions were
prepared in double-distilled water.
Vanadium(V) solutions were prepared by dissolving ammonium metavanadate in
the calculated quantity of sulfuric acid solution with vigorous shaking; the following
reaction occurs:
123

Oxidation of formaldehyde by pentavalent vanadium
Fig. 1 Representative first-order plot to evaluate the pseudo first-order rate constant (kobs) for
V(V) oxidation of formaldehyde in SDS catalyzed path in aqueous medium. [V(V)]T = 0.001 mol dm-3
[HCHO]T = 0.01 mol dm-3, [SDS]T = 9 mM, [H2SO4] = 3.0 mol dm-3, Temp. = 313 K
,
2H2SO4 þ NH4VO3 ! VOþ2 þ 2HSO4ꢂ þ NH4þ þ H2O
ð1Þ
Kinetic measurements
Our kinetic studies were always performed under pseudo first-order conditions by
using a known large excess of formaldehyde over ammonium metavanadate V(V).
The progress of the reaction was monitored by estimating the decay of V(V) from
UV–visible spectroscopic measurements (Shimadzu model 3600). The reaction was
initiated by adding the required amount of vanadium stock solution to the reaction
mixture. All kinetic runs were conducted in a thermostatic bath at 313 0.1 K. For
all solutions the UV-absorption maximum (k) was at 285 nm [46]. Pseudo first-
order rate constants (kobs) were obtained from the slopes of plots of -ln(A285
)
against time (t) (Fig. 1). Half life (t1/2) values were calculated directly by use of the
relationship t1/2 = ln2/kobs Under our experimental conditions the surfactants were
kinetically stable toward the redox activity of V(V). Hence, the rate of
decomposition of the surfactants by V(V) was not taken into consideration to
correct the observed rate constants. We repeated each experiment at least twice.
Method for CMC determination
Conductivity measurements
Conductivity was measured by use of a direct reading conductometer (Eutech
Instrument model PCD 6500) of cell constant 1 cm-1. All surfactant (SDS) solutions
were prepared in doubly-distilled water with a specific conductivity (j) of
1–3 lS cm-1. A pre-calibrated (by use of KCl solutions) conductivity cell was used
123

P. Sar et al.
Table 1 CMC and aggregate size (in aqueous medium) of the surfactants used to catalyse the oxidation
reaction
Surfactant
Type
Method
CMC (obs) (mM)
CMC (mM)a
Dh,Micelle (nm)b
SDS
Anionic
Anionic
Non-ionic
Conductometry
Fluorimetry
8.02
8.22
0.231
8.2
353
SDS
8.2
353
TX-100
Fluorimetry
0.24
12.68
Dh determined by DLS measurement
a
Reported CMC value
b
Hydrodynamic diameter
to obtain a specific conductance at an appropriate concentration. The cell, containing
solutions, was immersed in a water bath with temperature control variation of 0.1 K.
Trials were repeated three times to determine reproducibility. From the plot of
specific conductivity (j) vs. [SDS] we calculated the CMC of anionic SDS. For TX-
100, conductivity measurements are not possible because of its non-ionic character.
Fluorimetric measurements
Fluorescence measurements, at 313 K, were obtained by use of a fluorescence
spectrophotometer (Perkin Elmer model LS55) with excitation and emission slit
widths of 15.0 nm and 2.5 nm, respectively, and a scan speed of 500 nm min-1. For
SDS, excitation was at 493 nm and emission was recorded in the range
315–650 nm; for TX-100, excitation was at 250 nm and emission was recorded
in the range 280–500 nm. The concentration of the Congo red solution was fixed at
2 lM in different surfactant solutions. Typical experimental observations are
presented in Figs. S1 and S2. The CMC of the surfactants, calculated from the
breaks in plots of fluorescence intensity against [surfactant], are presented in
Table 1. The values are in agreement with those reported in the literature.
Scanning electron microscopy (SEM)
SEM is one of the most well known techniques used for determination of the surface
morphology of aggregates and the distribution of surfactant aggregates. For SEM
analysis, a drop of the sample solution was placed on foil paper, air dried, coated
with gold, and observed by field-emission scanning electron microscopy (Jeol JSM-
6700).
DLS measurements
To determine micelle and vesicle size and morphology, DLS characterization was
performed with a Malvern Zetasizer Nano ZS-90 (IACS, India) instrument with a
capillary cell. Size measurements were performed at an angle of 90°, in triplicate.
DLS analyses the velocity distribution of particle movement by measuring dynamic
123

Oxidation of formaldehyde by pentavalent vanadium
fluctuations of light-scattering intensity caused by the Brownian motion of the
particle. This technique yields a hydrodynamic radius to be determined via the
Stokes–Einstein equation from the aforementioned measurements:
kBT
Rh ¼
;
6pgoD
where D is the measured diffusion coefficient, kB is the Boltzmann constant, T is the
absolute temperature, and go is the solvent viscosity. From this equation, we can
also obtain values of the hydrodynamic diameter (Dh = 2Rh). DLS experiments
enable overall measurement of the particle perpendicular to the light source at that
moment. The light-scattering study was performed with SDS at a concentration
above the CMC value in 3.0 mol dm-3 H2SO4.
Product analysis and test for free radicals
Among all methods of separation, fractional distillation is one of the best for
separation of one component from a mixture. In this work the oxidation product was
efficiently isolated by fractional distillation. The 1H NMR spectrum of the separated
product was recorded at room temperature, in CDCl3 as solvent, by use of a Bruker
Ascend instrument operating at 400 MHz.
We confirmed the possibility of formation of free radical in our oxidation process
by use of acrylonitrile. When acrylonitrile was added to the reaction mixture a thick
semisolid was obtained, owing to polymerization initiated by free radicals [12].
Results and discussion
Colour change of the reaction mixture
It was stated above that vanadium has four common oxidation states (?5, ?4, ?3,
and ?2). The chemistry of vanadium(V) in aqueous solution is complex, and
depends on solution pH [47]. Here we used 3.0 mol dm-3 sulfuric acid (moderately
?
-
acidic medium) in which vanadium present as the VO2 ion (the VO3 ion in the
ammonium salt is converted to VO2? by acid in accordance with Eq. 1). Sadoc et al.
[48] reported a detailed computational study on the structure and stability of a
variety of vanadates and polyvanadates as a function of pH and vanadium
concentration.
The colour of ammonium metavanadate, (in which vanadium is present as V5?),
is yellow. During oxidation of formaldehyde the vanadium is reduced to V4?
(vanadyl ion, VO2?), which is blue in colour. The colours of the different reaction
mixtures are listed in Table 2. It was found that in the absence of the surfactant the
colour of the reaction mixture changed after several days whereas this particular
reaction was complete in a few minutes in the presence of the surfactant. It can,
therefore, be stated that transformation of V5? to V4? or formaldehyde to formic
acid is very slow in aqueous medium.
123

P. Sar et al.
Table 2 Colour changes during the oxidation process
Solution mixture
Colour of reaction
mixture
Oxidation state
of vanadium
Only NH4VO3 in H2SO4
Yellow
Yellow
Blue
V(?5)
V(?5)
V(?4)
V(?4)
Formaldehyde ? H? ? NH4VO3 (initial)
Formaldehyde ? H? ? NH4VO3 (after several days)
Formaldehyde ? H? ?NH4VO3 ? SDS (initial)
Blue
O
+
S
−
O
Na
O
O
SDS
H
O
O
9–10
TX-100
Scheme 1 Structures of SDS and TX-100
Characterization of the surfactants
The tendency of a surfactant is to form micelles, at the CMC, depends on the
hydrophobic character of alkyl tail of the surfactant [49]. When a surfactant forms
micelles, the aqueous surfactant solution behaves as a micro-heterogeneous system
[50].
From available surfactants we selected SDS (anionic) and TX-100 (non-ionic)
(Scheme 1) for our investigations. These surfactants are frequently studied and
commercially available with the same quality. The CMC of SDS was obtained by
two methods: conductance (Fig. 2a) and fluorescence (Fig. 2b), on the basis of
changes in specific conductance or fluorescence intensity as a function of [SDS].
The CMC value of TX-100 was obtained from the plot (Fig. 3) of fluorescence
intensity vs. [TX-100]. For SDS a sharp decrease in the fluorescence intensity of
Congo red occurs during micellization. For TX-100 (unlike SDS) the fluorescence
intensity initially decreases rapidly, in a linear fashion, with increasing [TX-100].
When micellization starts, the fluorescence intensity decreases less with increasing
[TX-100]. From the point of intersection of the two straight lines of different slope
we obtained the CMC of TX-100 (Fig. 3). The methods used for determination of
the CMC for both SDS and TX-100 are listed in Table 2. The CMC were obtained
as described elsewhere [50, 51].
123

Oxidation of formaldehyde by pentavalent vanadium
(a)
36
30
24
18
12
6
8.22 mM
0
6
8
10
12
[SDS] mM
T
(b)
Fig. 2 a Plot of specific conductance, j vs. concentration of SDS in aqueous solution at 313 K. b Plot of
fluorescence intensity vs. concentration of SDS in aqueous solution at 313 K
SDS and TX-100 both form micelles or vesicle-like aggregates under our
reaction conditions; the sizes of the aggregates were determined from DLS
measurements. From the size-distribution plots (Fig. 4) it was found that maximum
123

P. Sar et al.
Fig. 3 Plot of fluorescence intensity of Congo red as a function of TX-100 concentration at
kem = 250 nm at 313 K
Fig. 4 Size distributions of
36
30
24
18
12
6
micellar aggregates under our
reaction conditions (from DLS).
Inverted triangles, SDS at
9 mM; triangles, TX-100 at
0.28 mM. Under our reaction
conditions the maximum
number (intensity) of SDS
aggregates occurs at 353 nm
whereas for TX-100 it is at
12.68 nm only
0
0
100
200
300
400
500
600
Size (nm)
aggregate size for SDS (353 nm) was larger than for TX-100 (12.68 nm) (Table 1).
The micelle and/or vesicle-like aggregates of nanometre dimensions behave as
containers in which a variety of reactions occur. The nanometre-sized micelles
increase the local concentrations of the reactants around their surfaces, leading to
accelerated transformation.
Kinetics
Table 3 shows the pseudo first-order rate constants and half lives (t1/2) for the
oxidation in the absence and presence of surfactant. In aqueous solution the pseudo
123

Oxidation of formaldehyde by pentavalent vanadium
Table 3 Comparison of rate
constants in the presence and
absence of surfactants
Surfactant
–
103 9 micellar
106 9 kobs (s-1
)
t1/2 (h)
catalyst (mol dm-3
)
–
0.305 0.03
68.3 0.05
77.3 0.02
87.5 0.04
92.5 0.01
122.8 0.03
141.3 0.05
144.6 0.06
146.3 0.02
155.3 0.01
9.05 0.04
8.58 0.05
8.17 0.01
7.74 0.02
7.18 0.03
4.98 0.02
4.78 0.04
4.71 0.05
631.28
2.82
4.0
5.0
2.49
6.0
2.20
7.0
2.08
SDS
8.0
1.56
9.0
1.36
10.0
11.0
12.0
0.15
0.18
0.20
0.22
0.24
0.25
0.27
0.30
1.33
1.31
1.24
21.27
22.44
23.57
24.87
26.82
38.66
40.28
40.88
TX-100
[V(V)]T = 0.001 mol dm-3
,
[HCHO]T = 0.01 mol dm-3
,
[H2SO4] = 3.0 mol dm-3
Temp. = 313 K
,
first-order rate constant was kobs * 0.305 9 10-6 s-1 and the half life was
631.28 h. In the presence of SDS and TX-100 the rate constant is higher than in
absence of surfactant.
Effect of surfactants
We monitored the catalytic effects of the surfactants on the rate of oxidation of
formaldehyde. This study revealed the surfactant acts as a catalyst above its CMC but
not below its CMC. Initially, we determined the pseudo first-order rate constant for
oxidation of 0.001 mol dm-3 formaldehyde by 0.0001 mol dm-3 vanadium(V) in
aqueous medium of 3.0 mol dm-3 H2SO4 at 313 K. From the change in absorbance of
sequential spectra (Fig. 5) we determined the observed rate constant, for uncatalysed
oxidation, kobs * 0.305 9 10-6 s-1. Subsequently, reactions were performed in
presence of the surfactants (SDS and TX-100). Figure 6 shows sequential absorption
spectra for oxidation in the presence of SDS at regular time intervals. From Table 3 it
is apparent that the SDS-catalysed reaction proceeds 500 times faster than the
uncatalysed reaction and that TX-100-catalysed oxidation is 20 times faster than the
uncatalysed reaction. Therefore, we can say that non-functional micellar catalysts
SDS and TX-100 had a substantial effect in rate enhancement. The mechanisms of
catalysis are described below.
123

P. Sar et al.
Fig. 5 Sequential absorption spectra at regular time intervals (24 h) during the uncatalysed reaction.
[formaldehyde]T = 1 9 10-2 mol dm-3, [V(V)]T = 1 9 10-3 mol dm-3, [H2SO4] = 3.0 mol dm-3
Temp. = 313 K
,
Fig. 6 Sequential absorption spectra at regular time intervals (5 min) during the course of the catalysed
reaction. [formaldehyde]T = 1 9 10-2 mol dm-3
,
[V(V)]T = 1 9 10-3 mol dm-3
,
[H2SO4]T =
3.0 mol dm-3, [SDS]T = 10 9 10-3 mol dm-3, Temp. = 313 K
Catalytic effect of SDS
The effect of SDS on the rate of formaldehyde oxidation was studied at 313 K in the
concentration range [SDS]T = 4–12 mM and at constant H? ion, NH4VO3
(0.001 mol dm-3), and substrate (0.01 mol dm-3) concentrations. The results
presented in Table 3 show that rate constant of the oxidation reaction increases with
increasing concentration of SDS up to 7 mM. Above that concentration the rate of
oxidation does not increase further (i.e. is almost constant). The first observation is
123

Oxidation of formaldehyde by pentavalent vanadium
Fig. 7 Variation of the observed pseudo first-order rate constant (at k285 nm) with the concentration of
SDS at [H2SO4] 3.0 mol dm-3, Temp = 313 K. The abrupt change (increase) in the rate constant is
indicative of the CMC of SDS. The curve is the fit of the data to a simplistic sigmoidal equation
because of the formation of premicellar aggregates [52, 53] in the system (below the
CMC). From Fig. 7 and Table 3 it is apparent that kobs increased rapidly
from [SDS] = 7 mM (92.5 9 106 s-1) to [SDS] = 9 mM (141.3 9 106 s-1). This
observation gives essential information that during aggregation or micellization
(the experimented CMC of SDS is 8.02 mM) the rate of oxidation increases
much more quickly. The catalytic activity of SDS is high in this oxidation process,
because it increases the rate of this very slow oxidation reaction approximately
500-fold.
Catalytic effect of TX-100
Fascinating behaviour was observed for the non-ionic TX-100 in this oxidation
process. From Table 3 it is apparent that TX-100 enhanced the rate of oxidation
compared with the uncatalysed reaction. The rate of oxidation decreased gently with
increasing concentration of TX-100, however; then, at a specific concentration, it
suddenly decreased (Fig. 8). After this specific concentration the pseudo first-order
rate constant was not affected by increasing the concentration of TX-100. This
unusual behaviour is explained below.
It has been reported that the shape and size of micellar aggregates are affected by
surfactant concentration [54]. There is a gradual decrease in aggregate size (i.e.,
hydrodynamic diameter) with increasing surfactant concentration, because of high
electrostatic repulsion between the micelles [55]. As the size of the micelle
decreases with increasing concentration of TX-100, the number of reactive oxidant
species in close contact with micelle decreases, with the result that the oxidation
process slows down. The plot of rate constant versus concentration of TX-100
123

P. Sar et al.
Fig. 8 Variation of observed
pseudo first-order rate constant
(at k285 nm) with the
concentration of TX-100 at
[H2SO4] 3.0 mol dm-3
,
Temp. = 313 K. The abrupt
changes (decrease) in the rate
constant value inform us about
the CMC of TX-100
(Fig. 8) gives an idea about the reaction mechanism, and the micellization point of
TX-100.
SEM
In this work it was found that surfactants accelerate the slow oxidation reaction
whether surfactant is anionic or non-ionic. To explain this surprising behaviour we
assumed the substrate formaldehyde is encapsulated by the aggregate (or micelle).
This was later verified by UV–visible spectrophotometric analysis. Rispens and
Engberts [42] reported that a micellar solution acts as a micro-heterogeneous
system and reaction in the presence of micelle involves three steps: binding of the
substrate(s) with the micelle leads to chemical transformation at the micellar
surface and the end product is released from micelle. To investigate this we used
SEM. Detailed SEM study revealed that the substrate, formaldehyde, interacts with
the micelles formed by the corresponding surfactant. Figure 9a shows large uni-
lamellar vesicle type aggregates of SDS (10 mM) in 3.0 mol dm-3 H2SO4. This
informs us that vesicle-type aggregates are formed and are stable for a long period.
This result is in good agreement with the literature [33]. Figure 9b shows the
deformation of the structure of the micelle because of interaction with formal-
dehyde (0.01 mol dm-3) present in the micellar solution. Figure 9c shows the
deformation of the structure of micelle because of interaction with formaldehyde
(0.01 mol dm-3) and with ammonium metavanadate (0.001 mol dm-3) in the
micellar solution.
Comparing the three SEM pictures (Fig. 9a, b, c) we can say that the micelles
formed in 3.0 mol dm-3 H2SO4 are quite stable and interact with the substrate in
presence of oxidant species. Because the rate of oxidation increases, the micelle
used as catalyst must be interacting not only with the substrate but also with the
oxidizing species, vanadium(V). These pictures are in agreement with current views
on micelle formation.
123

Oxidation of formaldehyde by pentavalent vanadium
Fig. 9 Scanning electron microscopy images of SDS solution in H2SO4. a SDS only; b SDS in the
presence of formaldehyde; c SDS in the presence of formaldehyde and ammonium metavanadate. The
micellar aggregates change as
vanadium(V) species
a result of interaction among aggregates, formaldehyde, and
For SDS, the rate enhancement is because of the strong catalytic activity of the SDS
micelles. It has been reported that many reactions in micellar medium occurs either
inside the core of the micelle or in the ‘‘Stern layer’’ of the micelle; different factors
are responsible for accelerating or retarding reactions which occur in the micellar
stern layer region [56–59]. Here the oxidation reaction between formaldehyde and
vanadium(V) (oxidant) occurs in the micellar media when they meet at the micelle–
water interface (i.e., the Stern layer, which is less polar than water); i.e. the vanadium
complex is brought closer to the organic substrate by the surfactant head [60]. To
explain the micellar catalysis process, first we ensure that during the oxidation there is
only one type of surfactant aggregate and assume that the reactants do not
substantially change the structure of the aggregates. The results obtained by UV
analysis and SEM are very much in agreement with our assumptions. Multiple
surfactant aggregates are also apparent from optical microscopy (Fig. S5).
The explanation for enhancement of the rate of oxidation by SDS micellar media
is favourable electrostatic interaction between the anionic sulfate head-group and
the positive reactive species V(OH)3HSO4?, which results in the higher rate of
oxidation in the anionic micelle (SDS) than in the non-ionic micelles (TX-100) [21].
123

P. Sar et al.
Fig. 10 Variation of particle size distribution of SDS (9 mM) micelles and/or vesicles in 3.0 mol dm-3
H2SO4 determined by dynamic light scattering (DLS) measurement. The average particle sizes are
(squares) 353 nm, (circles) 401 nm, and (triangles) 456 nm
From the kinetic measurements (Fig. 8 for TX-100), it is apparent that the rate of
oxidation decreases with increasing [TX-100] for the entire concentration range and
that at the particular region of aggregate formation (CMC of TX-100 is
0.23 mM obtained from fluorescence measurement) the rate diminishes dramati-
cally. The explanation for the reaction in TX-100 medium is that the vanadium
complex probably cannot enter into the stern layer of the TX-100 micelle, although
the substrate formaldehyde reaches that layer.
DLS
The particle-size distributions obtained by DLS measurement are presented in
Fig. 10. The areas of the peaks correspond to the product of concentration and
particle mass. Assuming a common structure, the fraction with large diameter is
much less than 20 % of the total mass of the sample. From Fig. 10 it is apparent that
for SDS the size or diameter of the micelle changes when formaldehyde and
vanadium(V) are added to the micellar medium. The plot informs us about changes
in the micellar diameter when the reactants are present in the reaction medium. This
change in the size or diameter of the micelle helps us to state that an interaction
occurs between the substrate (formaldehyde) and vanadium(V) with the ‘‘Stern
layer’’ of the micelle. Table 1 shows the sizes of SDS and TX-100 micelles are
comparable. It has been found that under our reaction conditions the SDS micelle
(353 nm) is larger than the TX-100 micelle (12.68 nm). As a result, several oxidant
and substrate species can interact with the outer surface of aggregates formed by
SDS but not with those formed by TX-100. This is why the rate of oxidation in TX-
100 micellar medium is not as high as in SDS micellar medium.
123

Oxidation of formaldehyde by pentavalent vanadium
Table 4 Kinetic micellar effect, kmax/kw, for the oxidation reaction in the presence of surfactants in
3.0 mol dm-3 H2SO4 at 313 K
Surfactant
kmax/kw
SDS
509.18
29.67
TX-100
kmax pseudo first-order rate constant at the maximum of the surfactant profiles, kw pseudo first-order rate
constant in the absence of surfactant
In the absence of micelle or surfactant, the substrate (formaldehyde) and oxidant
(pentavalent vanadium) species are not very close to each other. The presence of the
micellar medium effectively brings them together, thus increasing the rate of the
reaction. The SDS micelle is anionic, therefore positive vanadium(V) species
interact with this and with formaldehyde molecules residing in the ‘‘Stern layer’’ of
the micelle.
Comparison of SDS and TX-100 catalysis
Rate constants for oxidation of formaldehyde by ammonium metavanadate, i.e.,
vanadium(V), in the presence of SDS and TX-100 at different concentrations are
presented in Table 3. The kinetic micellar effect is the ratio of the maximum pseudo
first-order rate constant in the presence of surfactant (kmax) to that in the absence of
surfactant (kw). The kinetic micellar effect for both SDS and TX-100 were
calculated by use of the data of Table 3. Table 4 shows the kmax/kw values obtained;
the kinetic micellar effect is much larger for SDS than for TX-100.
Another important factor which may affect the superior rate acceleration
behaviour of SDS over TX-100 is the hydrophilic–lipophilic balance (HLB). For a
single chain surfactant, HLB is the combined effect of the hydrophilic head group
and the hydrophobic tail, and ranges from 0 to 40. According to the literature the
HLB for SDS and TX-100 are 40 and 13.5, respectively. Detergents with higher
HLB are most useful for solubilisation of additives and/or substrates than those of
lower HLB [61]. Therefore, in this study we can say that SDS solubilises
formaldehyde better than TX-100. A higher HLB value of a surfactant corresponds
to greater hydrophobicity of the surfactant and, for this reason, there is a tendency of
monomers to assemble into a more lamellar aggregate. In this work SDS formed a
uni-lamellar type aggregate. Greater solubilisation and formation of large uni-
lamellar type aggregates are two important factors resulting in better catalysis by
SDS than TX-100.
Aggregation of surfactant monomers to form micelles results in catalysis of a
particular reaction. In a micellar catalysis reaction, there are different explanations
of rate enhancement [41]. The observed acceleration or retardation of the oxidation
rate has been explained on the basis of partitioning of the reactants in the micellar
and aqueous phase (Scheme 2).
The key factor in micellar catalysis is the location of the reacting species in the
relatively small volume of the micelles compared with the bulk solution. As a result
123

P. Sar et al.
kW
V(OH)3HSO4
Ks
Product
Product
Formaldehyde
Ks
W
M
W
M
kM
Formaldehyde
V(OH)3HSO4
Scheme 2 Partitioning of the reactive species and substrate between aqueous and micellar phases
Na+
Gouy-Chapman
Na+
Na+
Na+
Layer
Na+
Na+
Na+
Oδ−
5+
V
5
Vꢀ
Na+
+
δ
Na+
C
δ−
O
H
O
O
5+
V
O
−
δ+
C
H
H
Na+
S
OH
V
Na+
OH
HO
H
OH
5+
Electrostatic
Interaction
V
Na+
5+
V
Na+
Na+
H
Na+
Na+
Na+
H
Cδ+
H
5+
V
Na+
δ+
Oδ−
O
O
C
H
O
Na+
5+
V
δ−
O
Na+
5+
S
Na+
HO
HO
V
V
Na+
OH
Na+
OH
Stern Layer
Na+
Na+
Scheme 3 Interaction among the surfactant head group in the SDS micelle, formaldehyde, and the active
oxidant species
of this, a large increase in the effective concentration of reactants in the micellar
core results in rate enhancement.
A different observation from DLS measurement, is that the larger size of SDS
(353 nm) compared with TX-100 (12.68 nm). Larger aggregates have the ability to
increase the local concentration of reactant species (here vanadium species and
formaldehyde) much more than the smaller aggregates. Therefore, DLS data entirely
support the higher catalytic activity of SDS compared with TX-100.
123

Oxidation of formaldehyde by pentavalent vanadium
OH
HO
OH
O
O
O
+
H3O+
VO2
HSO4
+
+
OH
S
V
A
δ
−
OH
O
O
H
+
δ
C
V
C
A
H
OH
OSO2OH
H
H
1:1 Intermediate
Complex (C*)
O
H
C
fast
V(IV)
1:1 Intermediate
complex (C*)
+
Free radical
OH
Scheme 4 Vanadium(V) oxidation of formaldehyde
TX-100 increased the rate of the slow formaldehyde oxidation reaction, but not
as much as SDS, because of the negative charge on the head-group of SDS micelle,
which interacts strongly with the positive oxidant species. As a result the positive
oxidant species (vanadium ion) moves closer to the substrate (formaldehyde). All
these factors indicate that the reactive oxidant species and substrate are preferen-
tially partitioned at the SDS micelle–water interface, and to a larger extent than for
TX-100. Oxidation of formaldehyde by vanadium(V) in the SDS micelle is depicted
in Scheme 3.
Reaction scheme for formaldehyde oxidation
Vanadium(V) is known to form protonated species in acidic medium. The mechanism
of oxidation of formaldehyde is much affected by the different vanadium(V) species,
viz., VO2?, (VO.H2OꢁSO4)?, V(OSO3H)23?, (VO2ꢁH2OꢁH2SO4)?, V(OH)3HSO4
,
?
VO2ꢁHSO4, VO2(H2O)?4 , and VO3?, present in the sulfuric acid [62]. We propose
?
generation of positively charged active vanadium species by reaction of VO2 and
H2SO4:
123

P. Sar et al.
VOþ2 þ H3Oþ þ HSO4ꢂ ꢀ V(OH)3HSOþ4
.
?
After generation of V(OH)3HSO4 the formaldehyde undergoes nucleophilic
substitution. Reorganization of the positively charged substituted complex then
occurs via participation of a proton from the substrate [47]. Nikolov and Yatsimirskii
[47] reported that during the course of the oxidation reaction the vanadium complex
approaches the substrate, forming a 1:1 intermediate complex between substrate and
vanadium(V) (apparent from changes in the absorption intensity, Fig. S3). The
intermediate complex probably produces formic acid via a free radical mechanistic
pathway (formation of a thick semisolid precipitate on addition of acrylonitrile after
storage of the reaction mixture 72 days). After 72 days, the product was separated by
1
fractional distillation then a H NMR spectrum was recorded (supporting informa-
tion, Fig. S4). A plausible mechanism [62] of this reaction is given in Scheme 4.
Conclusions
Summary of this work
All the above experimental observations and discussion of this kinetic and
mechanistic study are summarized below.
(1) Oxidation of formaldehyde to industrially important formic acid by ammo-
nium metavanadate (V(V)) is a time-consuming reaction. The reaction is
accelerated by addition of surfactants.
(2) Two surfactants, SDS and TX-100, were used for micro-heterogeneous
catalysis of the reaction.
(3) The anionic micelle of SDS is an excellent catalyst compared with the non-
ionic micelle TX-100.
(4) SEM and DLS measurements of the micellar solutions under these reaction
conditions facilitate explanation of the catalytic effect of the SDS and TX-100
micelles on the reaction.
(5) A mechanism of the oxidation is proposed. It is supported by a variety of
experimental results and reported literature.
Before this work many research workers had reported that micellar catalysis
occurs via interaction between the micellar surface and the reagent (oxidant and
substrate) without providing any evidence. To the best of our knowledge, our
observations are the first experimental evidence of the interaction of reagent
(oxidant and substrate) with the micellar Stern-layer region.
With this system we achieved a simple, rapid, economical low cost, and
environmentally friendly method of oxidation to produce formic acid.
Acknowledgments We are grateful to UGC and CSIR, New Delhi, for financial support. The authors
would also like to recognize Mr Tathagata Mondal, at the Polymer Science Unit, and Mr Avijit Mondal,
of Advanced Materials, IACS, Kolkata, India, for their kind help providing SEM and DLS measurements.
123

Oxidation of formaldehyde by pentavalent vanadium
References
1. T. Hirao, Chem. Rev. 97, 2707–2724 (1997)
2. J.S. Littler, Tetrahedron 27, 81–91 (1971)
3. V. Jagannadham, Am. J. Chem. 2, 57–82 (2012)
4. R. Saha, R. Nandi, B. Saha, J. Coord. Chem. 64, 1782–1806 (2011)
5. B. Saha, C. Orvig, Coord. Chem. Rev. 254, 2959–2972 (2010)
6. P.C. Nagajyoti, K.D. Lee, T.V.M. Sreekanth, Environ. Chem. Lett. 8, 199–216 (2010)
7. A.S. Onkar, P.N. Naik, S.D. Gunagi, S.T. Nandibewoor, S.A. Chimatadar, Indian J. Chem. 51A,
1574–1579 (2012)
8. P.C. Wilkins, M.D. Johnson, A.A. Holder, D.C. Crans, Inorg. Chem. 45, 1471–1479 (2006)
9. M.A. Malik, F. Nabi, Z. Khan, J. Dispers. Sci. Technol. 29, 1396–1400 (2008)
10. H.K. Okoro, E.O. Odebunmi, Sci. Res. Essays 5, 2588–2594 (2010)
11. B. Saha, S. Sarkar, K.M. Chowdhury, Int. J. Chem. Kinet. 40, 282–286 (2008)
12. V.K. Sharma, K. Sharma, P.S. Tiwari, D. Khare, Int. J. Pharm. Life Sci. 2, 1223–1225 (2011)
13. Z. Khan, P.S.S. Babu, Kabir-ud-Din, Carbohydr. Res. 339, 133–140 (2004)
14. E.O. Odebunmi, A.S. Ogunlaja, S.O. Owalude, Orbital Electron. J. Chem. 2, 127–139 (2010)
15. S.T. Oyama, J. Phys. Chem. 94, 5029–5033 (1990)
16. R.M. Clementin, K. Takashima, Quim. Nova 16, 529–533 (1993)
17. S.K. Ghosh, A. Ghosh, R. Saha, K. Mukherjee, A. Basu, B. Saha, Tenside Surfactants Deterg. 49,
296–299 (2012)
18. S.K. Ghosh, A. Basu, K.K. Paul, B. Saha, Mol. Phys. 107, 615–619 (2009)
19. B. Saha, K.M. Chowdhury, J. Mandal, J. Solut. Chem. 37, 1321–1328 (2008)
20. B. Saha, Inorg. React. Mech. 6, 287–291 (2008)
21. K. Hanson, R.T. Baker, J.C. Gordon, B.L. Scott, L.A. Silks, D.L. Thorn, J. Am. Chem. Soc. 132,
17804–17816 (2010)
22. P.L. Timmanagoudar, G.A. Hiremath, S.T. Nandibewoor, J. Chem. Soc. Dalton Trans. 3623–3626
(1995)
23. A. Qiao, V.N. Kalevaru, J. Radnik, A.S. Kumar, N. Lingaiah, P.S. Sai Prasad, A. Martin, Catal.
Commun. 30, 45–50 (2013)
24. S. Wesselbaum, U. Hintermair, W. Leitner, Angew. Chem. Int. Ed. 51, 8585–8588 (2012)
25. D. Sodhi, M.A. Abraham, J.C. Summers, J. Air Waste Manag. Assoc. 40, 352–356 (1990)
26. J.Q. Torres, S. Royer, J.P. Bellat, J.M. Giraudon, J.F. Lamonier, ChemSusChem 6, 578–592 (2013)
27. L.R. Smith, Formic acid and formaldehyde destruction in waste streams, US Patent 1997, 5606107
28. T.J. Kemp, W.A. Waters, Proc. R. Soc. Lond. A Math. Phys. Sci. 274, 480–499 (1963)
29. A. Ghosh, R. Saha, P. Sar, B. Saha, J. Mol. Liq. 186, 122–130 (2013)
30. E.A. Batista, T. Iwasita, Langmuir 22, 7912–7916 (2006)
31. A.S. Narang, D. Delmarre, D. Gao, Int. J. Pharm. 345, 9–25 (2007)
32. C.O. Rangel-Yagui, A. Pessoa-Jr, L.C. Tavares, J. Pharm. Pharm. Sci. 8, 147–163 (2005)
33. J.H. Fendler, Annu. Rev. Phys. Chem. 35, 137–157 (1984)
34. J.B.F.N. Engberts, M. J. Blandamer, Chem. Commun. 1701–1708 (2001)
35. S. Otto, J.B.F.N. Engberts, Org. Biomol. Chem. 1, 2809–2820 (2003)
36. Z. Lin, R.M. Hill, H.T. Davis, L.E. Scriven, Y. Talmon, Langmuir 10, 1008–1011 (1994)
37. T. Shimanouchi, H. Umakoshi, R. Kuboi, Langmuir 25, 4835–4840 (2009)
38. G. Astray, A. Cid, J.A. Manso, J.C. Mejuto, O. Moldes, J. Morales, Int. J. Chem. Kinet. 43, 402–408
(2011)
39. C.A. Bunton, Catal. Rev. Sci. Eng. 20, 1–56 (1979)
40. E.H. Cordes, R.B. Dunlap, Acc. Chem. Res. 2, 329–337 (1969)
41. T. Dwars, E. Paetzold, G. Oehme, Angew. Chem. Int. Ed. 44, 7174–7199 (2005)
42. T. Rispens, J.B.F.N. Engberts, J. Org. Chem. 67, 7369–7377 (2002)
43. S.R.K. Minkler, B.H. Lipshutz, D.N. Krause, Angew. Chem. Int. Ed. 123, 7966–7969 (2011)
44. A. Ghosh, R. Saha, K. Mukherjee, S.K. Ghosh, S.S. Bhattacharyya, S. Laskar, B. Saha, Int. J. Chem.
Kinet. 45, 175–186 (2013)
45. A. Ghosh, R. Saha, K. Mukherjee, S.K. Ghosh, B. Saha, Spectrochim. Acta A Mol. Biomol. Spec-
trosc. 109, 55–67 (2013)
46. L. Newman, K.P. Qinlan, J. Am. Chem. Soc. 81, 547–549 (1959)
47. G.S. Nikolov, K.B. Yatsimirskii, Theor. Exp. Chem. 5, 549–560 (1969)
123

P. Sar et al.
48. A. Sadoc, S. Messaoudi, E. Furet, R. Gautier, E.L. Fur, L. Polles, J.Y. Pivan, Inorg. Chem. 46,
4835–4843 (2007)
49. D.W. Roberts, J. Costello, QSAR Comb. Sci. 22, 220–225 (2003)
50. A. Dominguez, A. Fernandez, N. Gonzalez, E. Iglesias, L.J. Montenegro, J. Chem. Educ. 74,
1227–1231 (1997)
51. S.K. Hait, S.P. Moulik, J. Surfactants Deterg. 4, 303–309 (2001)
52. P. Somasundaran, K.P. Ananthapadmanabhan, I.B. Ivanov, J. Colloid Interface Sci. 99, 128–135
(1984)
53. E. Iglesias, Int. J. Chem. Kinet. 44, 668–679 (2012)
54. V.K. Aswal, P.S. Goyal, Phys. Rev. E 61, 2947–2953 (2000)
55. M. Pisarcik, F. Devinsky, E. Svajdlenka, Colloids Surf. A 119, 115–122 (1996)
56. L. Garcia-Rio, P. Herves, J.R. Leis, J.C. Mejuto, P. Rodriguez-Dafonte, J. Phys. Org. Chem. 17,
1067–1072 (2004)
57. L. Onel, N.J. Buurma, J. Phys. Chem. B 115, 13199–13211 (2011)
58. P.C. Rey, C.C. Lago, P. Herves, J. Phys. Chem. B 114, 14004–14011 (2010)
59. L.S. Romsted, Langmuir 23, 414–424 (2007)
60. L. Kumar, T. Mahajan, D.D. Agarwal, Ind. Eng. Chem. Res. 51, 2227–2234 (2012)
61. B.W. Berger, Biophys. J. 89, 452–464 (2005)
62. A.P. Mishra, R. Khan, R.R. Pandey, Indian J. Chem. 48A, 1228–1234 (2009)
123