Kinetics of the molybdate and tungstate catalyzed oxidation of iodide by hydrogen peroxide

Article abstract of DOI:10.1016/S0020-1693(99)00079-1

Several studies of the Mo(VI) and W(VI) catalyzed reaction of H₂O₂ with various substrates have been documented previously. The experiments described herein attempt to settle some apparent contradictions in these reports. Specificall

Full text of DOI:10.1016/S0020-1693(99)00079-1

Inorganica Chimica Acta 288 (1999) 229–232
Note
Kinetics of the molybdate and tungstate catalyzed oxidation of
iodide by hydrogen peroxide
Christine L. Copper, Edward Koubek *
United States Na6al Academy, Annapolis, MD 21402, USA
Received 27 November 1998; accepted 4 February 1999
Abstract
Several studies of the Mo(VI) and W(VI) catalyzed reaction of H2O2 with various substrates have been documented previously.
The experiments described herein attempt to settle some apparent contradictions in these reports. Specifically, the oxidation of I−
with H2O2 in the presence of molybdate or tungstate ion in an acid solution was studied. The rate law for this system was found
to be ratei=kcat[MoO42−][I−]+kuncat[H2O2][I−] with values of kcat=3.87 M−1 s−1 and kuncat of 9.5×10−3 M−1 s−1 (at
22.8°C). The values of DH" for the molybdate catalyzed and uncatalyzed reactions were found to be 39.4 and 56.8 kJ mol−1
,
respectively. Similarly, the value of kcat was found to be 32.7 M−1 s−1 and DH" was found to be 41.1 kJ mol−1 (at 21.8°C) for
the same system catalyzed with tungstate rather than molybdate. © 1999 Elsevier Science S.A. All rights reserved.
Keywords: Catalysis; Iodide oxidation; Hydrogen peroxide; Molybdate; Tungstate
1. Introduction
H2MoO4+H2O2 ?H2MoO4(H2O2)
(rapid, K1)
(1)
(2)
(3)
Several studies of the Mo(VI) and W(VI) catalyzed
reaction of H2O2 with various substrates have been
reported [1–8]. One of the first detailed studies was that
of the oxidation of I− with H2O2 in the presence of
molybdate ion in an acid solution. This study by Garcia
and Lara [1] found that, in 0.08 M H2SO4 solution, the
rate law for the reaction is:
H2MoO4(H2O2)+H2O2 ?H2MoO4(H2O2)2
(rapid, K2)
H2MoO4(H2O2)2+I− ?products
(rate determining, k1)
s
with k1=3.3×102 M−1 −1. No activation parameters
were given. Studies by Thompson and co-workers [6]
have also shown that the formation of the proposed
peroxo complexes have large formation constants.
Recently, it has been reported [7] that the reaction
between H2O2 and an excess of I− in the presence of
molybdate in an acid solution follows the rate law:
rate=k[MoO4 ]
2− 0.5[H2O2]0.5[I−]1.2
Later, Smith and Kilford [2] studied the same reac-
tion with I− and H+ in excess over H2O2. They found
that under these conditions the reaction was neither
first nor second order in H2O2. A slight dependence on
H+ was found in the range 0.025–0.20 M. They con-
cluded that their results could best be explained by the
sequence:
rate=k2[Mo(VI)][H2O2]
with the formation of either a mono or diperoxo com-
plex being the rate determining step which is followed
by the rapid oxidation of I−.
We have decided to investigate this apparent contra-
diction by employing an iodide clock technique to
measure the true initial rate of reaction. We have
* Corresponding author. Tel.: +1-410-293 6634; fax: +1-410-293
2218.
0020-1693/99/$ - see front matter © 1999 Elsevier Science S.A. All rights reserved.
PII: S0020-1693(99)00079-1

230
C.L. Copper, E. Koubek / Inorganica Chimica Acta 288 (1999) 229–232
Table 1
Volumes of reagents (ml added to 25.0 ml buffered Na2S2O3 stock solution) and reaction times from varying [I−] and [H2O2]
Flask number
0.060 M KI
Distilled H2O
0.040 M H2O2
2.0×10−3 M Na2MoO4
Reaction time (s) at 22.8°C
1a
1b
2a
2b
3a
3b
4a
4b
5a
5b
25
25
25
25
25
25
15
15
10
10
29
30
19
20
0
0
9
10
14
15
20
20
30
30
50
50
50
50
50
50
1.0
0.0
1.0
0.0
1.0
0.0
1.0
0.0
1.0
0.0
55.8
113.7
45.0
78.6
33.1
47.3
56.1
79.9
83.0
118.0
chosen to work in an acetate buffered (pH 4.5) medium
where there is little interference by the acid promoted,
uncatalyzed (by MoO42−) reaction [9,10].
2.2. Timed reactions
2.2.1. Order in H2O2 and I−
To determine the order in H2O2, 25 ml of the
Na2S2O3, buffered, starch stock solution was trans-
ferred to a 250 ml Erlenmeyer flask followed by a
constant amount of Na2MoO4 and KI solution. Dis-
tilled water was added as indicated in Table 1. To begin
each reaction, varying amounts of H2O2 solution were
added, quickly mixed and the time noted. At the first
appearance of the blue starch-I2 complex, the time was
again noted. From this time and the change in thiosul-
fate concentration (from 2.5×10−4 to 0 M), the over-
all initial rate of each reaction was determined. By
2. Experimental
The reaction studied is a variation of the familiar
iodine clock reaction. The reaction is:
H2O2+2I− +2H+ “I2+2H2O
(4)
For each reaction mixture studied, the initial reaction
rate was determined by adding starch and a small,
known quantity of sodium thiosulfate to the iodide ion
solution before adding hydrogen peroxide. Upon mix-
ing the hydrogen peroxide with the iodide ion solution,
the molecular iodine initially produced Eq. (4) is imme-
diately consumed by the thiosulfate according to the
reaction:
Table 2
Catalyzed ratea dependence upon H2O2 and I− at pH 4.5, 22.8°C,
[MoO42−]=2.0×10−5
M
[H2O2]i
(M)
[I−
(M)
]
10−6 Catalytic rate
i
(M−1 s−1
)
I2+2S2O3 “2I− +S4O6
2−
(5)
2−
0.008
0.012
0.020
0.020
0.020
0.015
0.015
0.015
0.009
0.006
1.14
1.19
1.14
0.67
0.44
However, after the small, known quantity of thiosul-
fate is exhausted, the iodine produced will complex with
the starch resulting in a blue color in the reaction
solution. The interval between final reagent mixing and
the first appearance of a blue color throughout the
solution provides the time data for the mixture.
a Total rate-uncatalyzed rate=catalyzed rate.
2.1. Stock solutions
Since the rate of reaction of hydrogen peroxide with
iodide ion is greatly influenced by small amounts of
contaminants, care must be taken to insure the cleanli-
ness of all glassware.
A buffer-starch stock solution of Na2S2O3 was used
in all reactions. It was made by combining 25 ml of
0.01 M Na2S2O3, 25 ml of buffer (1 M HC2H3O2 and 1
M NaC2H3O2, pH 4.5), 25 ml of a 2% starch solution,
and 175 ml of distilled water.
2−
Fig. 1. Dependence of initial rate on MoO4
constant [H2O2] (=0.1155 M) and 22.8°C.
concentration at

C.L. Copper, E. Koubek / Inorganica Chimica Acta 288 (1999) 229–232
231
subtracting the uncatalyzed rates (those obtained from
the b flasks) from the overall rates (those obtained from
the a flasks), one obtains the initial catalyzed rates of
reaction. These initial rates for the Mo(VI) catalyzed
reactions, along with initial concentrations are given in
Table 2.
2.2.2. Order in Mo(VI) and W(VI)
The catalyzed reactions were run as before except that
2−
varying amounts of MoO4
(or WO42−) were added
2−
(Table 3). In the case of the rapid WO4
catalyzed
reactions, it was found that it was best to add the
Na2WO4 solution to the H2O2 solution before mixing.
The above processes were repeated for the catalyzed
and uncatalyzed reactions at a low temperature. To
maintain the low temperature, the flasks containing all
reagents were kept in a constant temperature bath prior
to use. Once the iodine clock reactions were begun, the
flasks were still kept at a constant, low temperature. No
attempt was made to maintain conditions of constant
ionic strength as changes in ionic strength are reported
to have no effect upon the rate of reaction [2,7].
3. Results and discussion
Our results (given in Tables 2 and 3) show that the rate
of I− oxidation at pH 4.5 in the presence of small
2−
catalyst quantities of MoO4
rate law:
can be expressed by the
ratei=kcat[MoO42−][I−]+kuncat[H2O2][I−]
Plots of the initial rate/[I−]i versus [MoO4 2−] at
constant [H2O2]i (see Fig. 1) were linear (r2=0.9996)
with a slope=kcat=3.87 M−1 s−1 and intercept=
1.1×10−4 at 22.8°C. Dividing the intercept by [H2O2]i
gives a value for kuncat of 9.5×10−3 M−1 s−1 which is
in excellent agreement with the value in the literature
[9]. Data obtained at 0.3°C also proved linear (r2=
0.9997) with a slope=kcat=1.03 M−1 s−1 and an
intercept=1.64×10−5
[H2O2]i gives a value for kuncat=1.42×10−3 M−1 s−1
. Dividing the intercept by
.
These results yield a DH" of 39.4 kJ mol−1 for the
molybdate catalyzed reaction and a DH" for the uncat-
alyzed reaction of 56.8 kJ mol−1. While the former
result is in poor agreement with the value reported by
Karunakaran and Muthukumaran [7] of 20.6 kJ
mol−1, the later is in good agreement with the value
reported (53.8 kJ mol−1) by Liebhafsky and Moham-
mud [9]. Similar results were obtained with the W(VI)
catalyzed reaction between I− and H2O2 (Table 3). kcat
at 21.8°C was 32.7 M−1 s−1 (r2=0.9970) while at
0.4°C kcat was 8.8 M−1 s−1 (r2=0.9970). This yields a
DH" for the catalyzed reaction of 41.1 kJ mol−1. It
should be noted that the Mo(VI) catalyzed reaction has
a slightly lower activation energy even though the
W(VI) is a superior catalyst. Thus the kinetic advantage

232
C.L. Copper, E. Koubek / Inorganica Chimica Acta 288 (1999) 229–232
of W(VI) has over Mo(VI) must be due to a more
favorable activation entropy. This result has been ob-
served previously on Mo(VI) and W(VI) catalyzed reac-
tions of H2O2 [3]. We are unable to offer a suitable
explanation for this result. It should be noted that the
kuncat for the uncatalyzed reactions obtained from this
W(VI) plot is again in excellent agreement with the
values in the literature [9].
One item to be noted is that our kcat for the Mo(VI)
catalyzed reaction is smaller than that reported by
Smith and Kilford [2]. This may be due to the differ-
ences in the pH and the reaction:
uncatalyzed pathway proceeds at a faster rate than the
catalyzed pathway. For example, it can be noted that in
flask 5 (Table 1) the catalyzed rate is 0.45×10−6 M−1
s
−1 while the uncatalyzed rate is 1.05×10−6 M−1 s−1
.
This is a direct result of the small quantity of catalyst
used.
One last item to be addressed is the rate law and
DH" reported by Karunakaran and Muthukumaran
[7]. As mentioned previously, they reported
rate=k[Mo(VI)][H2O2] with DH" =20.6 kJ mol−1
We believe the reason for the difference in their
results and those being reported here, is due to the
different initial conditions. They used a large excess of
I− and small quantities of H2O2 and Mo(VI). Appar-
ently, under these conditions, the formation of the
active Mo(VI) species is rate limiting. This leads to a
rate law that is first order in H2O2 and zero order in
I−. Our conditions are [I−]:[H2O2]ꢀ[Mo(VI)]. This
favors the rapid formation of the catalytic species
MoO(O2)2+H2O?Mo(OH)(O2)2− +H+
At higher pH values the species Mo(OH)(O2)2−
should be favored. Ghirm and Thompson have shown
this species to be a much less reactive species than
MoO(O2)2 [3]. This lowering of the rate of Mo(VI)
peroxide catalyzed reaction at higher pH has also been
noted before [11].
Thus, our results are essentially in agreement with
the mechanism originally proposed by Smith and Kil-
ford, i.e. at pH 4.5:
(DH" =20.6 kJ mol−1) and makes step (8) (DH"
=
39.4 kJ mol−1) rate limiting.
We have also investigated the effect of Mo(VI) upon
the reaction of t-butyl hydroperoxide and peroxydisul-
fate ion with I− and have found no catalysis. This
result is consistent with the above proposed catalytic
complex ions with two h2 equatorial peroxo ligands
[12].
H2O2+MoO42− ?MoO(OH)(O2)− +H2O
(rapid, K1)
(6)
(7)
(8)
MoO(OH)(O2)− +H2O2 ?Mo(OH)(O2)2− +H2O
(rapid, K2)
Mo(OH)(O2)2− +I− “products
References
(rate determining)
[1] F. Garcia, I. Gomez-Lara, Rev. Soc. Quim. Mex. 13 (1969)
222A.
along with the uncatalyzed reaction
H2O2+I− “HOI+OH− (rate determining)
[2] R.H. Smith, J. Kilford, Int. J. Chem. Kinet. 8 (1976) 1.
[3] A.F. Ghirm, R.C. Thompson, Inorg. Chem. 27 (1988) 4766.
[4] A.O. Chong, K.B. Sharpless, J. Org. Chem. 42 (1977) 1587.
[5] S.E. Jacobson, D.A. Muccigrosso, F. Mares, J. Org. Chem. 44
(1979) 921.
(9)
HOI+I− “products
(fast)
(10)
[6] J.D. Lydm, L.M. Schwane, R.C. Thompson, Inorg. Chem. 26
(1987) 2606.
[7] C. Karunakaran, B. Muthukumaran, Transition Met. Chem. 20
(1995) 460.
[8] V. Conte, F. Di Furia, in: G. Strukul (Ed.), Catalytic Oxidations
with Hydrogen Peroxide as Oxidant, Ch. 7, Kluwer, Boston,
1992.
[9] H.A. Liebhafsky, A. Mohammad, J. Am. Chem. Soc. 55 (1933)
3977.
However, our rate law is somewhat less complex than
those reported earlier [1,2]. The reason for this appears
to be, that under the conditions we have chosen,
[H2O2]ꢀ[Mo(VI)], the equilibrium proposed in steps 1
and 2 always lie far to the right. This keeps the concen-
tration of the active catalytic species constant and equal
to [MoO42−]i. This greatly simplifies the rate law for
the catalytic path. This large excess of H2O2 does lead
to a large contribution to the overall rate by the
uncatalyzed pathway: so much so that at times, the
[10] C.L. Copper, E. Koubek, J. Chem. Educ. 75 (1998) 87.
[11] F. Zhao-Lun, X. Shu-Kun, Anal. Chim. Acta 145 (1983) 143.
[12] R. Stomberg, Acta. Chem. Scand., Ser. A 42 (1988) 284.