Determination and evaluation of the second-order rate constants for the oxidation of maleic and fumaric acid derivatives with alkaline potassium permanganate

Article abstract of DOI:10.3184/030823409X424867

The second-order rate constants for the oxidation of maleic (Z-butenedioc) and fumaric (E-butenedioc) acid derivatives by potassium permanganate at pH 13 were determined. These constants varied by a factor of 104. The results suggested that the rates of attack of the permanganate ion on the double bonds were largely governed by steric factors with electrostatic and electronic effects playing only a minor role.

Full text of DOI:10.3184/030823409X424867

1
88 RESEARCH PAPER
MARCH, 188-190
JOURNAL OF CHEMICAL RESEARCH 2009
Determination and evaluation of the second-order rate constants for the
oxidation of maleic and fumaric acid derivatives with alkaline potassium
permanganate
S. Rahimah Abdul Halim, Gilly M. French, Alexandra J.N. Hughes, C.A. Olivia Prankerd Smith
and D.E. Peter Hughes*
Westminster School, London SWIP 3PB, UK
The second-order rate constants for the oxidation of maleic {Z-butenedioc} and fumaric {E-butenedioc} acid
derivatives by potassium permanganate at pH 13 were determined. These constants varied by a factor of 104•
The results suggested that the rates of attack of the permanganate ion on the double bonds were largely governed
by steric factors with electrostatic and electronic effects playing only a minor role.
Keywords: butenedioic acid, alkaline permanganate, oxidation, tartaric acid
With dilute alkaline permanganate solutions maleic and
fumaric acids give meso and racemic tartaric acid respectively,
showing that the addition of the two hydroxyl groups is syn.
The oxidation of these acids (and other alkenes) by potassium
permanganate in acidic and neutral conditions has been
extensively studied,l-4 but there are many fewer data for
oxidation under alkaline conditions. In alkaline solution, it is
thought that there are two main steps in the reaction, leading
to the formation of the vic dioI.5
Other workers who used acid or neutral conditions2,4for the
oxidation expressed their rate in terms of -d[alkene]/dt. In
order to be able to compare rates under different conditions
of alkalinity, the rates in terms of -d[Mn04-]/dt rates were
converted into rates in terms of -d[alkene]/dt by dividing
by two.
The pseudo first-order rate constant for the decomposition
of permanganate (no alkene) in the presence of 0.5 mol dm-3
NaOH (an excess) was 0.0106 ± 0.0006 S-l. This value was
subtracted from the pseudo first-order rate constants in the
presence of alkene to give corrected rate constants.
Mn04-(aq) + -CH=CH-{aq) ---> Mn(Y) adduct
The rates were measured by using a spectrophotometer
set to 510 nm. At this wavelength the ratio of the absorbance
of Mn04-(aq) to the absorbance of MnOi-(aq) was ~6,
Mn(Y) adduct + 20H-(aq) + Mn04-(aq) --->
-CHOHCHOH-{aq) + 2MnOi-(aq)
a
maximum value. The absorbances of Mn04-(aq) and
Earlier work with E-but-2-enoic acidS had shown that the rate
was first order with respect to permanganate and first order
with respect to alkene if the reaction conditions were such
that the first step was much slower than the second. In this
work, the concentrations of the fumaric acid and maleic acid
derivatives were chosen so that this was the case.
Mn042-(aq) were found to be linear with concentration up
to a concentration of 5 x 10-4mol dm-3 which was the upper
concentration limit used; thus as the reaction took place there
was a linear relationship between [Mn04-(aq)] and absorbance.
In all experiments an excess of the alkene was used. Under
these conditions [Mn04-(aq)] decreased exponentially with
time showing a pseudo first-order relationship. A curve fitting
program was used to find the rate constant for the reaction.
The order with respect to maleate is shown in Table 1.
The constancy of the second-order rate constant confirms
that the reaction is first order in maleate.
The Mn042-(aq) ion was stable to disproportionation if
[0 H-(aq)] was greater than 1mol dm-3; at lower concentrations
disproportionation took place, although this reaction was
very slow at 0.5 mol dm-3 (pH ~13), which was the usual
concentration used in this work.
The order of the reaction with respect to hydroxide ion
is shown in Table 2, with [maleate] kept constant at
0.040 mol dm-3.
The ionic strength of the solution was kept at 2.0 mol dm-3
with sodium perchlorate. Because the rate of decomposition
of the Mn04- ion varied with [OH-], it was necessary to
High concentrations of OH-(aq) ions promoted the following
decomposition reaction.
These two competing effects were minimised by choosing
Table
1
Rate constants when [maleate] was varied
[OH-(aq)] = 0.5 mol dm-3. At this concentration, the rate of
disproportionation was very slow and could be neglected,
while the decomposition rate was slow but measurable.
The rate of decomposition was subtracted from the rates of
reaction using alkene to make allowance for this accompanying
reaction.
[Maleate]/
mol dm-
Corrected pseudo
first-order rate
constantsis-1
Second-order
rate constant
/ dm3 mol-1 S-1
3
0
.01
0.015
.02
Ionic strength
Temperature = 25.0°C.
0.060
0.085
0.117
±
±
±
0.001
0.003
0.006
3.0
2.9
±
±
0.1
0.2
0
2.9 ± 0.3
0.5 mol dm-3 . [OH-]
0.46-48 mol dm-3 .
Results and discussion
It was established that 1.0 mol of alkene reacted quantitatively
with 2.0 mol of permanganate. As the rate of the reaction
was followed by the disappearance of the purple Mn04- ion,
this gave the rate of the reaction in terms of -d[Mn04-]/dt.
The overall stoichiometric equation was 2Mn04-(aq) + =C=C=
Table
OH-]
2
Rate constants when [OH-] was varied.
[
Corrected pseudo first-order rate constant/s-1
0.5
1.0
1.5
2.0
0.049
0.042
0.044
0.048
(aq) + 20H-(aq) ---> =COHCOH = (aq) + 2MnOi-(aq).
*
Correspondent.E-mail:dephI5@hotmai1.com

JOURNAL OF CHEMICAL RESEARCH 2009 189
allow for this by subtracting a different constant from each
measured rate constant to give the corrected rate constant.
The [maleate] was 0.040 mol dm,3; thus the second-order rate
q~
~~
/110
'\
~
0
O
constant was 0.046/(0.040 x 2)
=
0.57 dm3 moP s'l, a value
- - M n - O
1
/6th of that determined when the ionic strength was 0.5 mol
/
\.
1
o
0
dm,3. This decrease in rate as the ionic strength was increased
was consistent with the attack of a negatively charged Mn04'
ion at a region of high electron density.
The second-order rate constants for a variety of maleic and
fumaric acid derivatives were determined by this method; each
determination was repeated several times in order to check for
consistency in the results. The results are given in Table 3.
The structure of the manganese complex formed is still
uncertain although some computational methods suggest,7,8
that 2 is preferred over 1.
The principal difference in these results at pH 13 to those
obtained by other workers in acid or neutral conditions2,4
is that their measured rates of reaction with maleic acid
derivatives were very much faster than those found here.
It may be that at a lower pH, some isomerisation to the more
stable fumaric derivative had taken place which would have
a big effect on the reaction rate. The only other work that has
been done at pH 13 was on the ethyl esters in 25% ethanol;
this showed that the fumaric ester reacted 50 times as fast as
Temperature
=
25°C. Ionic strength
=
0.50 mol dm,3.
The rate was independent of [OH'] but was kept within the
range 0.4-0.5 mol dm,3 to minimise side reactions.
The most obvious feature of these results is the slowness
of the rates with maleic acid derivatives compared to the
equivalent fumaric acid derivatives (about 200 times). It has
been suggested that the slow rates of reaction with maleic acid
derivatives are due to repulsion between the carboxyl groups as
they are forced closer together when the Sp2hybridised carbon
atoms change to sp3 hybridisation while the intermediate is
the maleic ester which is much nearer to the value obtained
9
here (200 times). This big difference in rates emphasised the
importance of the repulsion of the carboxyl groups in the
transition state of maleic acid derivatives.
The results suggest that the repulsion of the carboxy I
groups and the ester groups in the maleic acid derivatives is
similar, showing that the repulsion is largely steric, rather than
electrostatic, in nature. Substituents adjacent to the double
bond slowed down the rate of reaction, irrespective of whether
they were electron donating (methyl group) or electron
withdrawing (bromine atom).The reduction in rate was most
pronounced with the introduction of the large phenyl group
(3 times) and the second bromine atom (40 times).
being formed.
1
In alkaline conditions, this repulsion could
be either electrostatic or steric in origin. Such repulsion is
absent with fumaric acid and the rates of its derivatives will
be considered first.
The rates of reaction with fumaric acid derivatives and with
propenoate and E-but-2-enoate are very similar. If electronic
effects played a significant part in determining the rate, then
there would have been a marked difference between the
rate with E-but-2-enoate and with the rates with fumaric
acid derivatives which contain an extra carboxyl group.
This suggests that the initial attack showed no strong
electrophilic or nucleophilic activity. Such behaviour is
consistent with the initial formation of a cyclic organometallic
complex 1, as has been suggested when the reaction was
carried out at pH 6.86.3
This complex can either be directly converted into the
vic diol or could go via the cyclic diester 2; there is some
evidence that more than one oxygen is transferred from
the permanganate to the double bond, which would support
the formation of the cyclic diester.6
There were several advantages in carrying out the oxidation
at pH 13 rather than at a lower pH. The first was that the
degree of ionisation of the carboxyl groups was known
(approaching 100%); the second was that there was much less
likelihood of the oxidation being carried beyond the vic diol
stage; the third that there was less likelihood of isomerisation
of the maleic acid derivatives and fourthly measurement of
the absorption of a homogeneous solution was likely to show
more consistent results than measurement of the absorption of
a solution containing colloidal manganese dioxide.
Experimental
All the chemicals were obtained from Sigma-Aldrichand glass
distilled water was used for the solutions.The solutions of maleic
acid, bromomaleaicacid and dichloromaleicacid were made from
the anhydrides which minimised the possibility of fumaric acid
impurities. The esters and halogeno acids were made up an hour
before the rates were carried out and there was no measurable
hydrolysisduringthat time.Therewas alsono measurablehydrolysis
during the few minutes when they were mixed with the sodium
hydroxidein the cuvette.
Table
3
Values for second-order rate constants for some
maleic and fumaric acid derivatives
Alkene
Corrected second-order rate
constant/ dm3 mol-1 S-1
Maleic acid derivatives
Thestoichiometryofthepermanganate/fumarat er eactionwasfound
02CCH=CHC02
2.9 ± 0.3
2.6 ± 0.1
1.7 ± 0.05
by adding 10.0cm
3
of 0.0020mol dm-
3
potassiumpermanganateto
of 1.0mol dm-
02CCMe =CHC02
1
0.0em of 0.0010mol dm-
3
3
fumaricacidin 10cm
3
3
Me02CCH=CHC02Me
NaOH. Extrapermanganatewas then addeduntil the green solution
became tinged with purple. This correspondedto a stoichiometry
of 2.08 to 1. When allowancewas made for the decompositionof
permanganateas it was mixedwiththe sodiumhydroxide,this result
indicatedthat the reactionquantitativelygavethe diol.
The rate experiments were carried out in a Camspec M30I
spectrophotometer.The reaction vessel was a cuvette (path length
I cm), kept at 25°C by circulatingwater from a thermostatically
controlledbath.Theoutputwasfedto a PicoScopeADCI00analogue
to digital converterand hence to a computer.In a typical run, 200
resultswererecorded,eachonebeingthe averageof severalthousand
readings from the spectrophotometer.The resulting results were
analysedwith a curve fittingprogramwhich showedthat a good fit
0
0
2CCBr =CHC02
2CCPh =CHC02
1.8 ± 0.2
0.87 ± 0.03
0.070 ± 0.004
0.16 ± 0.01
-02CCBr =CBr02-
02CCCI =CCIC02
Fumaric acid derivatives
02CCH=CHC02
610 ± 30
480 ± 20
470 ± 20
630 ± 50
02CCMe =CHC02
Me02CCH=CHC02Me
Me02CCH=CHC02
Others
CH2=CHC02
MeCH=CHC02
560 ± 30
520 ± 20
(
to at leastfourhalf-lives)with an exponentialplot.Withmaleicacid
derivativesthe rate was slow enough for the reactantsto be mixed

1
90 JOURNAL OF CHEMICAL RESEARCH 2009
manually before starting tbe computer; the initial concentration ofthe
maleic acid derivatives varied between 0.0 I and 0.5 mol dm- which
was always in excess of the initial [Mn04-] which was 5 x 10-4mol
dm-3• Witb fumaric acid derivatives a modified technique was used;
Received 16 October 2008; accepted 28 December 2008
Paper 08/0234 doi: 10.3184/030823409X424867
Published online: 6 April 2009
3
1
.0 cm
3
oftbe 0.5 mol dm-
3
alkali and 1.0 cm
3
of 0.0010 mol dm-
3
References
alkene were mixed together in the cuvette, tbe computer started and
tben 100 flL of 0.0020 mol dm-3 Mn04- rapidly injected just below
tbe surface from a Hamilton syringe, fitted with a ptfe plunger. This
resulted in tborough mixing of the Mn04- ion in less tban 0.5 s and
allowed determination of rate constants with a halflife of ~ I s without
having to use a stopped-flow reactor. The initial concentrations oftbe
1
2
3
4
5
6
7
8
9
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3
and that oftbe alkene 2.0-
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2
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