A kinetic study of d-glucose oxidation by bromine in aqueous solutions

Article abstract of DOI:10.1016/j.carres.2006.04.035

The kinetics of the oxidation of d-glucose to d-gluconic acid by bromine in aqueous solution were studied using potentiometric techniques and theoretical considerations of complex bromine-bromide-pH equilibria. The pH has a strong influence on reaction rate. At pH < 8 the reaction is very slow, while in the pH range pH 8-9.5 the reaction is sufficiently fast and seems optimal for the reaction. The proposed active species at that pH region is hypobromous acid. At pH > 9.5, the reaction is further accelerated due to the formation of hypobromite. The proposed kinetics expression for gluconic acid formation, based on the determined kinetic parameters at pH 9.24, is of the form. dc(GA)/dt = 160c²(G)c₀(HOBr)c⁰(H⁺)c⁰(Br^-).

Full text of DOI:10.1016/j.carres.2006.04.035

Carbohydrate Research 341 (2006) 1779–1787
A kinetic study of D-glucose oxidation by bromine
in aqueous solutions
a
a
Branimir N. Grgur,a, Dragana L. Zugic, Milica M. Gvozdenovic
*
ˇ
´
´
b
´
and Tomislav Lj. Trisˇovic
aFaculty of Technology and Metallurgy, University of Belgrade, Karnegijeva 4, 11020 Belgrade, Serbia and Montenegro
bInstitute of Technical Science, Serbian Academy of Science and Arts, Knez Mihailova 35, 11020 Belgrade, Serbia and Montenegro
Received 30 January 2006; received in revised form 13 April 2006; accepted 19 April 2006
Available online 11 May 2006
Abstract—The kinetics of the oxidation of D-glucose to D-gluconic acid by bromine in aqueous solution were studied using poten-
tiometric techniques and theoretical considerations of complex bromine–bromide–pH equilibria. The pH has a strong influence on
reaction rate. At pH < 8 the reaction is very slow, while in the pH range pH 8–9.5 the reaction is sufficiently fast and seems optimal
for the reaction. The proposed active species at that pH region is hypobromous acid. At pH > 9.5, the reaction is further accelerated
due to the formation of hypobromite.
The proposed kinetics expression for gluconic acid formation, based on the determined kinetic parameters at pH 9.24, is of the form
dc(GA)/dt = 160c2(G)c0(HOBr)c0(Hþ)c0(Br)
ꢀ 2006 Elsevier Ltd. All rights reserved.
Keywords: Bromine; D-Glucose; D-Gluconic acid; Hypobromous acid; Oxidation; Indirect electroorganic synthesis
1. Introduction
This reaction is a well-known reaction found in organic
chemistry textbooks,6 but there is a lack of information
concerning the kinetics of this oxidation. The earliest
studies of electrochemical glucose oxidation to gluconic
acid were carried out by Isbell and co-workers.7,8 Bro-
mine ions in the presence of calcium carbonate were first
electrochemically oxidized to hypobromite, which chem-
ically oxidized glucose via a d-gluconolactone intermedi-
ate. Fink and Summers9 found that the optimum
electrolyte composition for electrolysis in the semi-
industrial production of calcium gluconate was 0.8 M
D-glucose and 0.17 M sodium bromide. Even though
such a process gives a high yield of product without
by-products, it has been replaced with enzymatic pro-
cesses10 with long reaction times and relatively low selec-
tivity or by oxidation with air in the presence of noble
metal catalysts, which can be easily deactivated over
time.11–13 The main reasons for giving up the electro-
chemical process were use of degradable carbon anodes
Indirect electroorganic synthesis is based on the electro-
chemical production of a halide active species in situ,
which reacts in solution with organic reactants. In general,
the reaction is very important due to a high-efficiency pro-
cess that is easy to carry out. Examples of these reactions
include the production of gluconic acid or calcium gluco-
nate by glucose oxidation1–3 or xylonic acid from xylose4
with electrochemically generated bromine, or indirect oxi-
dation of cyclohexanol to cyclohexanone with iodine.5
D-Glucose oxidation in aqueous solutions containing
bromine is shown in Eq. 1
d-glucose + Br2 + H2O ! d-gluconic acid + 2HBr
ð1Þ
*
Corresponding author. Tel.: +381 113303681; fax: +381
113370387; e-mail: BNGrgur@tmf.bg.ac.yu
0008-6215/$ - see front matter ꢀ 2006 Elsevier Ltd. All rights reserved.
doi:10.1016/j.carres.2006.04.035

1780
B. N. Grgur et al. / Carbohydrate Research 341 (2006) 1779–1787
and the high cost of electric power at that time. Today,
dimension-stable anodes (DSA) are widespread and the
cost of electric power is comparably less. One more
advantage of electrochemical processes is the possibility
of paired electroorganic synthesis, for example, the pro-
duction of D-gluconic acid and D-glucitol (sorbitol) at
the same time.2,3 Hence, consideration of electrochemi-
cal production of D-gluconic acid or D-gluconates is
again of interest, and a detailed kinetics study should
be the first step in processes optimization.
2. Experimental
Potentiometric titration was used to study the kinetics of
D-glucose oxidation to D-gluconic acid. Pt mesh served
as the indicator electrode, and a saturated calomel elec-
trode (Er = 0.243 V) served as the reference electrode.
All equilibrium potentials in this study are given versus
the standard hydrogen electrode. The potential measure-
ments were carried out with a PAR 273 potentiostat con-
trolled by a computer through a GPBI PC2A interface.
In order to determine the parameters of D-glucose oxi-
dation, solutions of bromine and D-glucose were ob-
tained by dissolution of stock solutions, using buffers
to maintain constant pH.
In aqueous solutions on an indicator electrode
(e.g., Pt), bromine in the presence of free electrons
establishes a bromine–bromide electrochemical equili-
brium:14
The bromine stock solution, concentration of 0.1 M in
0.2 M KBr, was obtained by dissolving KBr and calcu-
lated volume of the liquid bromine (q = 3.1 g cmꢀ3) in
doubly distilled water. This solution was protected from
photolytic degradation by storing in a dark, cold place.
The stock solution of D-glucose (2.50 M) was prepared
by dissolving solid, anhydrous a-D-glucose (p.a. E.
Merck) in doubly distilled water. This solution after
storage, due to mutarotation, can be treated as an ano-
meric mixture of a and b isomers.
The buffer solutions used were as follows: for the pH
interval of 7.50–9.24 a combination of 0.05 M sodium
tetraborate (borax) and 0.1 M HCl (Titrival); for the
pH interval of 9.24–12.4 a combination of 0.05 M
sodium tetraborate and 0.10 M NaOH, and for pH
4.01 and 12.9, 0.05 M potassium diphthalate and
0.10 M NaOH, respectively.
The relative volumes of the solutions of borax and
HCl or NaOH required for a total volume of 100 mL
at a given pH are shown in Figure 1.
Equilibrium potential measurements were performed
by adding bromine stock solutions (3–10 0.01 mL)
to the 100 mL of buffer solutions stirred by a magnetic
stirrer. After a stabile value of the equilibrium potential
had been established, D-glucose stock solution was
(Pt)Br2 + 2e ¡ 2Brꢀ
ð2Þ
or in the presence of bromide ions, a tribromide is
formed:
ðPtÞBr3ꢀ þ 2e ¡ 3Brꢀ
ð2aÞ
The equations describing the changes of equilibrium
potentials, Er, are defined as
2:303RT
2F
a2ðBrꢀÞ
aðBr2Þ
ErðBr2jBrꢀÞ ¼ ErhðBr2jBrꢀÞ ꢀ
Ehr ¼ 1:0874 V
log
ð3Þ
and
2:303RT
a3ðBrꢀÞ
log
ErðBr3ꢀjBrꢀÞ ¼ Ehr ðBr3ꢀjBrꢀÞ ꢀ
Ehr ¼ 1:053 V
ꢀ
2F
aðBr3 Þ
ð3aÞ
where Ehr refers to standard equilibrium potentials. After
rearrangements, these equations can be used to monitor
bromine or tribromide activities over time, where Er(t) is
the potential as a function of time:
a2ðBrꢀÞ
h
i
aðBr2Þ ¼
ð4Þ
h
ðtÞꢀE
ꢁ
ꢀ2F ½E
r
r
2:303RT
10
100
90
or
V=100-V (borax)
x
a3ðBrꢀÞ
aðBr3ꢀÞ ¼
ð4aÞ
0.1 M HCl
h
i
80
h
ðtÞꢀE
ꢁ
ꢀ2F ½E
r
r
0.1 M NaOH
2:303RT
10
70
60
50
40
(Note: in dilute aqueous solutions, activities can be
approximately replaced by concentrations.)
As a consequence, the kinetics of D-glucose oxidation
can be monitored in situ by measuring electrochemical
potential in time (potentiometry).
7.5
8.0
8.5
9.0
9.5
10.0
10.5
11.0
11.5
12.0
12.5
pH
In this paper, we present the kinetic studies of D-glu-
cose oxidation by bromine in aqueous solution, with
consideration of the various thermodynamic equilibria,
present in solution as a function of pH.
Figure 1. Required volume of 0.05 M sodium tetraborate for buffer
solution preparation of 100 mL combined with 0.10 M HCl or 0.10 M
NaOH.

B. N. Grgur et al. / Carbohydrate Research 341 (2006) 1779–1787
1781
added (0.2–3 0.01 mL), taking this time as t = 0. It
was assumed that addition of usually 3.2 mL to maxi-
mum 10 mL of stock solutions in 100 mL of a buffer
solution does not change pH significantly. All experi-
ments were performed at a temperature of 25 ꢁC. A
circulating constant-temperature bath maintained the
temperature of the electrolyte within 0.5 ꢁC. Tempera-
ture control was done by thermometer with an accuracy
of 0.1 ꢁC.
(c) The establishment of chemical equilibria and for-
mation of tribromide as well as hypobromous acid
and hypobromite at different pH of the solutions
should be considered.
In aqueous solutions, the dependence of bromine and
bromide ion concentration at the pH of the solution is
very complex. In order to determine which of the bro-
mine species is dominant at a specific pH, the following
thermodynamic analysis was carried out, based on the
known chemical reactions in this complex system.
In aqueous solutions of bromide ions and bromine the
following chemical equilibria are observed.15,16
3. Results and discussion
Figure 2a shows the time dependence of the equilibrium
potentialatthePtindicatorelectrode[(Pt)Br2(aq)jBrꢀ(aq)]
in solution containing 3 mM of total bromine and 0.20 M
KBr after addition of 75 mM of D-glucose. It can be seen
that the potential changes are very slow, over ꢂ3000 s
only ꢂ20 mV. By using Eqs. 4 and 4a, the changes of bro-
mine and tribromide concentrations over time can be
calculated. These are shown in the inset of Figure 2.
The following conclusions can be made based on the
data from Figure 2.
(1) Disproportionation with formation of hypobro-
mous acid and bromide ions:
K1h
Br2 þ 2H2O ¢ HOBr þ H3Oþ þ Brꢀ
ð5Þ
The standard chemical equilibrium constant for
the reaction, Kh1, is equal to its formal equilibrium
constant, K1, keeping in mind that Rm = 0:
aðHOBrÞaðH3OþÞaðBrꢀÞ
Kh1 ¼ K1 ¼
¼ 7:2 ꢃ 10ꢀ9
aðBr2Þa2ðH2OÞ
(a) The calculated initial concentration of free bromine
in the solution at t = 0 was ꢂ1 mM, which is
approximately three times smaller compared to
the concentration of totally added bromine
(3 mM). This is probably due to a rapid reaction
between bromine and bromide ions (Eq. 2a), with
formation of tribromide ions. The calculated initial
concentration for tribromide ions was ꢂ3 mM.
(b) The oxidation of D-glucose by bromine is a pseudo
first-order reaction with respect to bromine. This
can be determined from the linear dependences of
the logarithm of the bromine and tribromide con-
centrations as a function of time.
ð5aÞ
(2) Dissociation of hypobromous acid to hypobromite
and a proton:
K2h
HOBr þ H2O ¢ BrOꢀ þ H3Oþ
aðBrOꢀÞaðH3OþÞ
ð6Þ
Kh2 ¼ K2 ¼
¼ 2 ꢃ 10ꢀ9
ð6aÞ
aðHOBrÞaðH2OÞ
(3) Disproportionation with formation of hypobromite
and bromide ions:
K3h
Br2 þ 2OHꢀ ¢ BrOꢀ þ Brꢀ þ H2O
aðBrOꢀÞaðBrꢀÞaðH2OÞ
ð7Þ
1.05
a)
3 mM (Br2)T
75 mM D-glucose
0.2 M KBr
pH~6
Kh3 ¼ K3 ¼
¼ 2 ꢃ 108
aðBr2Þa2ðOHꢀÞ
1.04
1.03
1.02
1.01
1.00
ð7aÞ
10-2
10-3
Formation of tribromide ions:
K4h
Br-3|Br-
Br2 þ Brꢀ ¢ Brꢀ3
ð8Þ
Br2|Br-
aðBrꢀ3 Þ
Kh4 ¼
¼ 16:85 Mꢀ1
ð8aÞ
10-4
aðBr2ÞaðBrꢀÞ
0
1000
2000
3000
t / s
The mass-balance equation with respect to free bro-
mine, c(Br)s, in the solution can be used to determine
the concentration of the bromine species as a function
of pH (Note: for a detailed analysis, see Supplementary
data), with the approximation that activities could be
0
500
1000
1500
2000
2500
3000
t / s
Figure 2. Time dependence of equilibrium potential after addition of
D-glucose. Inset: Time dependences of bromine and tribromide
concentrations as calculated from Eqs. 4 and 4a.

1782
B. N. Grgur et al. / Carbohydrate Research 341 (2006) 1779–1787
110
100
90
80
70
60
50
40
30
20
10
0
replaced by concentrations and that water has unity
activity:
0.0030
0.0025
0.0020
0.0015
0.0010
0.0005
0.0000
3.0 mM (Br2)T
0.2 M KBr
cðBr2Þs ¼ cðBr2ÞT ꢀ cðBr3ꢀÞ ꢀ cðHOBrÞ
Br-3
-
BrO
ꢀ cðBrOꢀÞ ꢀ cðBrOꢀÞHOBr
ð9Þ
where c(Br2)T is the total analytical concentration of
bromine in the solution and c(BrOꢀ)HOBr is the concen-
tration of hypobromite ions, produced by hydrolysis of
hypobromous acid, given by
(Br2)s
HOBr
10
K2cðHOBrÞ
cðBrOꢀÞHOBr
¼
ð10Þ
cðH3OþÞ
Dividing Eq. 9 by c(Br)s, the following was obtained:
0
2
4
6
8
12
14
cðBr2ÞT cðBr3ꢀÞ cðHOBrÞ cðBrOꢀÞ
pH
1 ¼
ꢀ
ꢀ
ꢀ
cðBr2Þs cðBr2Þs
cðBr2Þs
cðBr2Þs
Figure 3. Concentration (left) and percentage (right) pH distribution
of bromine species in a solution of 3 mM of total bromine in 0.20 M
KBr.
cðBrOꢀÞHOBr
ꢀ
ð11Þ
cðBr2Þs
Finally, after introducing Eqs. 5a–8a and 10 in Eq. 11
and rearranging, the concentration of free bromine in
the solution as a function of pH was obtained:
As it can be seen in Figure 3, the pH of the solution has
great influence on the bromine species distribution,
which, on the other hand, could have great influence on
the kinetics of the D-glucose oxidation. At pH lower than
pH 6, the dominant species are free bromine and tribro-
mide ions, present at ꢂ23% and 77%, respectively. At
pH greater than pH 6, the concentrations of free bromine
and tribromide ions decrease, while the concentrations of
hypobromous acid and hypobromite ions increase. The
maximum concentration of hypobromous acid is realized
at pH 8.3 (ratio of ꢂ47 %), after which it decreases due to
dissociation into hypobromite and a proton.
ꢀ
K1
cðBr2Þs ¼ cðBr2ÞT 1 þ Kh4cðBrꢀÞ þ
cðBrꢀÞcðH3OþÞ
ꢁꢀ1
K3c2ðOHꢀÞ
K1K2
þ
þ
ð12Þ
cðBrꢀÞ
cðBrꢀÞc2ðH3OþÞ
By an analogous procedure, once the dependence of
the free bromine concentration on pH is known, it is
also possible to calculate the pH dependence of the con-
centration of all other species in the solution.
For hypobromous acid
The presence of different bromine species in the solu-
tion also leads to establishment of different electrochem-
ical equilibria that are characterized by their own
equilibrium potentials:
ꢂ
ꢃꢀ1
ꢂ
K2
cðBr2ÞT
cðBr2Þs
cðHOBrÞ ¼ ꢀcðBr2Þs 1 þ
1 ꢀ
cðH3OþÞ
ꢃ
K3c2ðOHꢀÞ
þKh4cðBrꢀÞ þ
ð13Þ
ð14Þ
ð15Þ
Br2 þ 2e¡2Brꢀ
ð16Þ
cðBrꢀÞ
2:303RT
2F
c2ðBrꢀÞ
cðBr2Þ
for total hypobromite ions
Er ¼ Ehr ðBr2jBrꢀÞ ꢀ
log
Ehr ¼ 1:0874 V
ꢂ
cðBr2Þ
cðBrOꢀÞ ¼ ꢀcðBr2Þ 1 ꢀ
þ Kh4cðBrꢀÞ
cðBr2ÞTs
s
ð17Þ
ð18Þ
ꢃ
K1
Brꢀ3 þ 2e¡3Brꢀ
þ
cðBrꢀÞcðH3OþÞ
2:303RT
2F
c3ðBrꢀÞ
cðBrꢀ3 Þ
Er ¼ Ehr ðBr3ꢀjBrꢀÞ ꢀ
log
Ehr ¼ 1:0503 V
and for tribromide ions
ꢂ
cðBr2Þ
K1
cðBrꢀ3 Þ ¼ ꢀcðBr2Þ 1 ꢀ
þ
cðBr2ÞTs
cðH3OþÞcðBrꢀÞ
s
ð19Þ
ð20Þ
ꢃ
K3c2ðOHꢀÞ
K1K2
HOBr þ H3Oþ þ 2e¡Brꢀ þ 2H2O
2:303RT
þ
þ
cðBrꢀÞ
cðBrꢀÞc2ðH3OþÞ
cðBrꢀÞ
Er ¼ Ehr ðHOBrjBrꢀÞ ꢀ
Ehr ¼ 1:33 V
log
Based on the analysis given above, the theoretical pH
dependence of concentration of all considered bromine
species in the solution is given in Figure 3. for a total bro-
mine concentration, c(Br2)T, of 3 mM in a solution of
0.20 M KBr.
cðHOBrÞcðH3OþÞ
2F
ð21Þ
ð22Þ
BrOꢀ þ H2O þ 2e¡Brꢀ þ 2OHꢀ

B. N. Grgur et al. / Carbohydrate Research 341 (2006) 1779–1787
1783
2:303RT
2F
cðBrꢀÞc2ðOHꢀÞ
Er ¼ Ehr ðBrOꢀjBrꢀÞ ꢀ
Ehr ¼ 0:761 V
log
3.0
10-1
3 mM (Br2)T
25 mM D-glucose
0.2 M KBr
cðBrOꢀÞ
2.5
2.0
1.5
1.0
0.5
0.0
ð23Þ
10-2
Formation of the D-gluconic acid from D-glucose can
be described by the general kinetic equation:
c(HOBr)
c(OBr- )
c(Br2)s
c(Br3-)
10-3
10-4
10-5
dcðGAÞ
¼ kcpðGÞcqðBrÞcrðHþÞcsðBrꢀÞ
ð24Þ
dt
where GA and G refer to D-gluconic acid or D-gluconate
and D-glucose, respectively. Br is the bromine species
participating in the reaction, k is the rate constant,
and p, q, r, and s refer to partial reaction orders of the
participants in the reaction.
4
6
8
10
12
14
pH
Figure 5. pH dependence of the slopes obtained from linear parts of
the curves in Figure 4 compared with theoretical distribution curves of
bromine species.
As already mentioned, the pH has great influence on
the bromine species distribution. Therefore, its influence
on the rate formation of D-gluconic acid was investigated
first. The results are given in Figure 4. The experimental
conditions had been chosen in such a way that the start-
ing concentrations of D-glucose and bromide ion were
scientifically greater than the total bromine concentra-
tion, in order to insure pseudo first-order conditions with
respect to bromine, at the beginning of the reaction.
As can be seen in Figure 5, the reaction practically
does not take place at pH 4.0 and 6.0, that is, only extre-
mely small changes of potential were observed in a short
period. Hence, it could be concluded that free bromine
and tribromide ions are only weakly active in the oxida-
tion of D-glucose. With increasing pH, the reaction
accelerates and in the range of pH 6–10 after introduc-
ing D-glucose into the solution, significant changes in
the equilibrium potential were observed. With further
increase of the pH, at pH P 10, the reaction further
accelerates, followed by a rapid drop of potential in less
than 5 s (Fig. 4). With the assumption of pseudo first-or-
der with respect to bromine, the slope (dEr/dt) should be
proportional to the pseudo first-order rate constants.
The actual values of the slope as a function of pH are
given in Figure 5 (bold line) and are compared to the
theoretical distribution of the bromine species, as estab-
lished earlier (Fig. 3).
Since the value of the slope is proportional to the rate
constant, the dependence of the logarithm of the slope
on the logarithm of the hydrogen-ion concentration,
that is, pH, pH = ꢀloga(H+), directly gives the reaction
order with respect to pH. It follows that in the pH
region between pH ꢂ6 and 8, the reaction order with
respect to H+ is about ꢀ1; in the region between pH ꢂ8
and 9.5, the order is ꢂ0; in the region between pH
ꢂ9.5 and 11, about ꢀ1, and at even greater pH again
around ꢂ0. If these results are compared with the distri-
bution curves of the bromine species as a function of pH
(Fig. 3), the rate of the reaction can be matched with a
particular bromine species. In the pH region between
pH ꢂ7 and ꢂ9.5, the dominant species should be hypo-
bromous acid, while at higher pH, the dominate species
could be hypobromite ion. All this illustrated the com-
plexity of the D-glucose oxidation with Br2, which re-
sults from the existence of various equilibria in the
bromine–bromide ion system.
1.10
1.05
3 mM (Br2)T
25 mM D-glucose
1.00
0.95
0.90
0.85
0.80
0.75
0.70
0.65
0.60
0.55
0.2 M KBr
pH
Considering that in acid and in weakly acidic solu-
tions, the D-glucose oxidation does not effectively take
place, the optimum conditions for the reaction would
be alkaline, that is, pH > 10. However, in alkaline solu-
tion, even at room temperature, the disproportionation
reaction of hypobromite ions to bromate and bromide
ions occurs according to:14
4.0 and 6.0
7.62
8.14
8.80
9.24
10.0
11.1
12.4
12.9
dEr/dt
3BrOꢀ ¡ BrOꢀ3 þ 2Brꢀ
ð25Þ
Although as shown in Figure 6, the bromate ion is a
relatively strong oxidant that can, in principle, oxidize
D-glucose, it is obvious from Eq. 25 that 2 mol of the
active BrOꢀ species is lost, which diminishes the overall
effectiveness of the reaction.
0
5
10
15
20
25
30
t / s
Figure 4. Time dependence of equilibrium potentials at different pH
values of the solution.

1784
B. N. Grgur et al. / Carbohydrate Research 341 (2006) 1779–1787
0.8
0.7
0.6
0.5
0.4
0.3
0.2
0.1
assumed that the active species in the reaction is hypo-
bromous acid due to the highest standard potential
equation 21, and that its concentration differs from the
concentration of total added bromine, the problem of
determining the hypobromous acid concentration arises.
For that reason, the actual concentration of hypobro-
mous acid was calculated using a rearranged form of
Eq. 21 (solved for c(HOBr) starting with the concentra-
tion c0(HOBr) at time t = 0). Again assuming a pseudo
first-order dependence, the following can be written:
25 mM D-glucose
0.2 M KBr
pH=12.88
3 mM (Br2)T
3 mM BrO-3
dcðGAÞ
dt
dcðHOBrÞ
dt
¼ ꢀ
¼ k0cðHOBrÞ
ð26Þ
where k0 is the pseudo first-order rate constant, k0 =
kcp(G)cs(Brꢀ).
Integrating Eq. 26 and introducing for t = 0,
c(HOBr) = c0(HOBr), one obtains
-10
0
10
20
30
t / s
cðHOBrÞ
c0ðHOBrÞ
k0
2:303
Figure 6. Comparison of time dependence of equilibrium potentials
after introduction of D-glucose into a solution containing 3 mM
bromine and 3 mM bromate ions.
log
¼ ꢀ
t
ð27Þ
Figure 8 shows the time dependence of the ratio of the
logarithms of actual and zero time concentrations of
hypobromous acid, calculated using the rearranged
Eq. 21. Through this procedure, a family of straight
lines with essentially identical slopes was obtained. This
is a confirmation of the pseudo first-order assumption.
Based on Eq. 27, the slope seen in Figure 8 is propor-
tional to the pseudo first-order rate constant, and the
average value of the rate constant is k0 = ꢀ2.303 ·
Hence, the optimal pH for D-glucose oxidation lies
between pH 8 and 9.5. Another advantage of operating
in weakly alkaline solution is the faster hydrolysis of
the initial reaction product (D-glucono-1.5-lactone) of
D-glucose to D-gluconic acid or at this pH to D-gluco-
nate. Fast hydrolysis of D-glucono-1.5-lactone in weak
alkaline media is completed in few minutes, while in acid
media this reaction is very slow.17
Figure 7 shows time dependences of equilibrium
potentials for different concentrations of total added
bromine at pH 9.24, while holding the concentrations
of D-glucose and bromide ion constant.
slope = 0.075 0.01 sꢀ1
.
On the other hand, considering Eq. 21 and a rear-
ranged Eq. 27:
2:303RT
2F
cðBrꢀÞ
Er ¼ Ehr ðHOBrjBrꢀÞ ꢀ
log
cðHOBrÞcðH3OþÞ
It can be seen that for the first 60 s, the time depen-
dence of the potential is essentially linear. Since it was
ð21Þ
1.02
0.2 M KBr
pH=9.24
25 mM D-glucose
0.2 M KBr
pH=9.24
25 mM D-glucose
1.00
100
c(Br2)T / mM
c(Br2)T / mM
3
3
0.98
5
7
5
7
-1
10
0.96
10
10
0.94
0.92
0.90
-2
slope = - 0.032+/-0.005 s-1
10
-3
10
0
20
40
60
80
100
0
20
40
60
80
100
t / s
t / s
Figure 7. Time dependence of equilibrium potential for different
concentrations of total bromine with constant concentration of
D-glucose and bromide ions.
Figure 8. Time dependence of the logarithm of the ratio of actual and
zero time concentration of hypobromous acid, for different total
bromine concentrations.

B. N. Grgur et al. / Carbohydrate Research 341 (2006) 1779–1787
1785
c0ðHOBrÞ
cðHOBrÞ
k0
2:303
log
¼
t
ð27aÞ
1.00
0.95
0.90
0.85
0.80
the equilibrium potential can be given as a function of
reaction time as
2:303RT
2F
cðBrꢀÞ
3 mM (Br2)T
0.2 M KBr
pH=9.24
Er ¼ Ehr ðHOBrjBrꢀÞ ꢀ
log
c0ðHOBrÞcðH3OþÞ
0 RT
2F
ꢀ k
t
ð28Þ
c(G) / mM
5
or at constant pH
12.5
25
50
75
0 RT
2F
Er ¼ const ꢀ k
t
ð29Þ
So, it is possible to determine the value of the rate
constant from the slope of the linear part of the equilib-
rium potential time dependence:
0
20
40
60
80
100
120
140
2F
RT
t / s
k0 ¼ ꢀ
slope
ð30Þ
Figure 9. Time dependence of measured equilibrium potentials for
The constant k0 is defined for constant values of bro-
mine and bromide ions concentration at zero time by
varying glucose concentration, according to
different D-glucose concentrations.
Table 1. The value of the slope determined for different D-glucose
concentrations (Fig. 9) for the constant concentrations of c(Br2)T
3 mM and calculated rate constant, k0, using Eq. 30
k0 ¼ kcsðBrꢀÞcpðGÞ ¼ k00cpðGÞ
ð31Þ
=
Considering Eq. 30, one obtains
c(Glucose) (M)
ꢀSlope (V sꢀ1
)
k0 (sꢀ1
)
RT
2F
k00cpðGÞ
ð32Þ
0.005
0.0125
0.025
0.05
2.97 · 10ꢀ5
2.01 · 10ꢀ4
0.00121
0.00235
0.0159
0.0951
0.30816
0.56721
ꢀslope ¼
and by taking the logarithmic form of a rearranged Eq.
32:
0.00389
0.075
0.00716
RT
2F
logðꢀslopeÞ ꢀ log
¼ log k00 þ p log cðGÞ
ð33Þ
102
k''= 160 s-1 M-1
a determination of the partial reaction order, p, for
D-glucose and the rate constant, k00, is possible.
In this manner, values of equilibrium potential as a
function of time were measured for different D-glucose
concentrations as shown in Figure 9.
As it can be seen in Figure 9, an initial linear depen-
dence of equilibrium potential on time is observed,
and the constant k0 can be calculated from the slope
using Eq. 30, as shown in Table 1.
3 mM Br2
0.2 M KBr
pH=9.24
101
100
10-1
10-2
10-3
Slope = 2.06
If Eqs. 31 and 33 are graphically presented, the reac-
tion order with respect to D-glucose can be calculated
from the slope, and from the intercept at a glucose con-
centration of 1.0 M, the value of the rate constant, k00, is
obtained. These results are given in Figures 10 and 11.
The reaction order determined for D-glucose by this
method is 2, and the calculated second-order rate con-
10-2
10-1
100
log [c(D-glycose) / M]
stant k00 is essentially the same, 160 Mꢀ1 sꢀ1
.
Figure 10. Graphical presentation of Eq. 31.
From Eq. 21, it is obvious that the value of the equi-
librium potential is also dependent on the bromide ion
concentration. Therefore, in order to complete the rate
equation, it is necessary to determine the reaction order
of bromide ion as well. In Figure 12 the equilibrium po-
tential change as a function of time for different bromide
ion concentrations is shown, while keeping the pH and
total bromine concentration constant.
From Figure 12 it can be seen that for reaction times
less than 20 s, the equilibrium potential value dropped

1786
B. N. Grgur et al. / Carbohydrate Research 341 (2006) 1779–1787
2.5
2.0
c(Br-) = 0.1 M
1 mM (Br2)T
6.0x10-4
3 mM (Br2)T
0.2 M KBr
pH=9.24
log k''= 2.2
1.5
0.2 M
1.0
0.5
4.0x10-4
0.0
pH=9.24
0.3 M
0.5 M
-0.5
-1.0
-1.5
-2.0
-2.5
-3.0
2.0x10-4
slope = 2.06
0.0
5
6
7
8
9
10
11
12
10-2
10-1
100
pH
log [c(D-glucose) / M]
Figure 13. Distribution of hypobromous acid concentration at con-
stant total bromine concentration of 1 mM and different bromide ion
concentrations and pH values.
Figure 11. Graphical presentation of Eq. 33.
ꢀ
kinetics of Br2 and Br3 ion hydrolysis. It is surprising
1 mM (Br2)T
75 mM D-glucose
pH 9.24
1.00
0.95
0.90
0.85
0.80
0.75
that the value of the reaction order of bromide ion is
zero, considering that the bromide ion has great influ-
ence on the distribution of the bromine species in the
solution. However, if the distribution curves for hypo-
bromous acid are drawn, for the constant value of the
bromide ion concentration, as shown in Figure 13, it is
evident that at pH values greater than 9, the bromide
ion concentration has no influence on the change of
hypobromous acid concentration. By this, the change
in reaction order of the bromide ion is explained. This
also confirms that the active species in the D-glucose oxi-
dation at the beginning of the reaction is really hypobro-
mous acid.
c(Br-) / M
0.1
0.2
0.3
0.5
Based on the parameters determined at pH 9.24, the
kinetic equation for D-glucose oxidation to D-gluconate
by Br2 is therefore
0
20
40
60
80
100
120
dcðGAÞ
¼ kcpðGÞcqðBrÞcrðHþÞcsðBrꢀÞ
ð24Þ
t / s
dt
Figure 12. Time dependences of measured equilibrium potentials at
constant total bromine concentration of 1 mM and different bromide
concentrations.
which can be rewritten as
dcðGAÞ
¼ 160c2ðGÞc0ðBrꢀÞc0ðHþÞc0ðHOBrÞ
¼ 160c2ðGÞc0ðHOBrÞ
dt
by ꢂ90 mV, equivalent to a decrease of the concentra-
tion of the active species by three orders of magnitude.
For short reaction times, the change in bromide ion con-
centration has essentially no influence on the reaction
rate. Therefore, it can be assumed that the reaction or-
der with respect to bromide ion is equal to zero. For
times longer than 20 s, the reaction rate decreased with
increase of bromide ion concentration, which means
that the reaction order is negative. However, for time
greater than 20 s, the concentration of the active species
could be neglected, and the reaction is determined by the
ð34Þ
where c0(HOBr) can be explicitly calculated by combin-
ing Eqs. 12 and 13.
The measurable equilibrium potential change as a
function of time can be obtained from Eq. 28 as
2:303RT
2F
cðBrꢀÞ
Er ¼ Ehr ðHOBrjBrꢀÞ ꢀ
log
c0ðHOBrÞcðH3OþÞ
RT
2F
ꢀ 160
c2ðGÞt
ð35Þ

B. N. Grgur et al. / Carbohydrate Research 341 (2006) 1779–1787
1787
4. Conclusions
References
1. Li, H.; Li, W.; Guo, Z.; Gu, D.; Cai, S.; Fujishima, A..
Collect. Czech. Chem. Commun. 1995, 60, 928–934.
2. Yu, J. C.; Boizer, M. M.; Nobe, K. J. Electrochem. Soc.
1988, 135, 1400–1406.
3. Pintauro, P. N.; Johnson, D. K.; Park, K.; Baizer, M. M.;
Nobe, K. J. Appl. Electrochem. 1984, 14, 209–220.
The kinetics of D-glucose oxidation to D-gluconic acid in
aqueous solution are investigated by potentiometric
technique. The reaction is strongly pH dependent, as
the pH affects the bromine and bromide ion equilibrium.
At pH 9.24, kinetic parameters for the reaction were
determined, assuming that the dominant active bromine
species is hypobromous acid.
´
´
ˇ ´
´
´
4. Jokic, A.; Ristic, N.; Jaksica, M.; Spasojevic, M.; Krstajic,
N. J. Appl. Electrochem. 1991, 21, 321–326.
5. Roquero, P.; Cognet, P.; Duverneuil, P.; Lacoste, G.;
Fabre, P. L. J. Appl. Electrochem. 1998, 28, 475–
482.
The kinetic parameters at this pH are as follows:
(1) The reaction order in hydrogen ion is zero.
6. Loudon, G. M. In Organic Chemistry; Oxford University
Press Inc.: New York, 2001; p 1328.
7. Isbell, H. S.; Frush, H. L. J. Res. Natl. Bur. Stand. 1931, 6,
1145.
8. Isbell, H. S.; Frush, H. L.; Butes, F. J. Ind. Eng. Chem.
1932, 24, 375–378.
(2) The reaction is pseudo first order in hypobromous
acid, and the rate constant is 0.075 sꢀ1
.
(3) The reaction is second order in D-glucose, and the
determined rate constant is 160 Mꢀ1 sꢀ1
.
(4) The reaction order in bromide ion is zero.
9. Fink, C. G.; Summers, D. B. Trans. Electrochem. Soc.
1938, 74, 625–629.
10. Riegel, E. R.; Bissinger, H. G. Riegel’s Handbook of
Industrial Chemistry; Kluwer Academic–Plenum Publi-
shers: New York, 2003; pp 992–994.
The kinetic equation for the formations of D-gluconic
acid formation is proposed in the form
dc(GA)/dt = 160c2(G)c0(HOBr)c0(Hþ)c0(Brꢀ)
11. Gangwal, V. R.; Schaaf, J.; Kuster, B. F. M.; Schouten,
J. C. Catal. Today 2004, 96, 223–234.
12. Onal, Y.; Schimpf, S.; Claus, P. J. Catal. 2004, 223, 122–
¨
133.
Acknowledgments
13. Abbadi, A.; Bekkum, H. J. Mol. Catal. A: Chem. 1998, 97,
111–118.
14. Mussini, T.; Loghi, P. The Halogens. In Standard Poten-
tials in Aqueous Solution; Bard, A. J., Parsons, R., Jordan,
J., Eds.; Marcell Dekker: New York, 1985; pp 70–
83.
This work was financially supported by the Ministry
of Science and Environmental protection, Republic of
Serbia, under Contract No. 142044.
15. Popov, A. I. In Halogen Chemistry; Guttman, V., Ed.;
Academic Press: New York, 1967; Vol. 1, pp 225–
227.
16. Downs, A. J.; Adams, C. J. In Comprehensive Inorganic
Chemistry; Trotman-Dickenson, A. F., Ed.; Pergamon
Press: New York, 1973; pp 1191–1194.
Supplementary data
Supplementary material associated with this article can
be found, in the online version, at doi:10.1016/j.carres.
2006.04.035.
17. Dijken, J. P.; Tuijl, A.; Luttik, M. A. H.; Middelhoven,
W. J.; Pronk, J. T. J. Bacteriol. 2002, 184, 672–678.