Silver(I) and Copper(II) Catalysis for Oxidation of Histidine by Cerium(IV) in Acid Medium: A Comparative Kinetic Study

Article abstract of DOI:10.1002/kin.21063

The catalytic effect of silver(I) and copper(II) ions on the oxidation of histidine by cerium(IV) in aqueous sulfuric acid solutions was studied spectrophotometrically at a constant ionic strength of 3.0 mol dm^?3 and at 25°C. In both uncatalyzed and metal ions-catalyzed paths, the reactions exhibited first-order kinetics with respect to [Ce(IV)] and [catalyst], and fractional first-order dependences with respect to [His] and [H⁺]. The oxidation rates increased as the ionic strength and dielectric constant of the reactions media increased. The catalytic efficiency of Ag(I) was higher than that of Cu(II). Plausible mechanistic schemes for both uncatalyzed and catalyzed reactions were proposed, and the rate laws associated with the suggested mechanisms were derived. In both cases, the final oxidation products of histidine were identified as 2-imidazole acetaldehyde, ammonium ion, and carbon dioxide. The activation parameters associated with the second-order rate constants were evaluated.

Full text of DOI:10.1002/kin.21063

Silver(I) and Copper(II)
Catalysis for Oxidation of
Histidine by Cerium(IV) in
Acid Medium: A
Comparative Kinetic Study
AHMED FAWZY,1,2 ISMAIL I. ALTHAGAFI,1 HATEM M. ALTASS1
1Chemistry Department, Faculty of Applied Sciences, Umm Al-Qura University, Makkah, Saudi Arabia
2Chemistry Department, Faculty of Science, Assiut University, Assiut, Egypt
Received 1 September 2016; revised 19 November 2016; accepted 1 December 2016
DOI 10.1002/kin.21063
Published online 24 January 2017 in Wiley Online Library (wileyonlinelibrary.com).
ABSTRACT: The catalytic effect of silver(I) and copper(II) ions on the oxidation of histidine by
cerium(IV) in aqueous sulfuric acid solutions was studied spectrophotometrically at a constant
ionic strength of 3.0 mol dm−3 and at 25°C. In both uncatalyzed and metal ions-catalyzed
paths, the reactions exhibited first-order kinetics with respect to [Ce(IV)] and [catalyst], and
fractional first-order dependences with respect to [His] and [H+]. The oxidation rates increased
as the ionic strength and dielectric constant of the reactions media increased. The catalytic
efficiency of Ag(I) was higher than that of Cu(II). Plausible mechanistic schemes for both
uncatalyzed and catalyzed reactions were proposed, and the rate laws associated with the
suggested mechanisms were derived. In both cases, the final oxidation products of histidine
were identified as 2-imidazole acetaldehyde, ammonium ion, and carbon dioxide.ꢀTChe activation
parameters associated with the second-order rate constants were evaluated.
Periodicals, Inc. Int J Chem Kinet 49: 143–156, 2017
2017 Wiley
ity. Cerium(IV) is less stable in aqueous nitric and
perchloric acid solutions [9–12]. It was rarely used in
perchloric acid medium due to the presence of dimers
and polymers of it in such medium [11,12]. Identifica-
tion of the kinetically active Ce(IV) species [2,3] is a
problem during studies of the kinetics and mechanisms
of cerium(IV) oxidation in aqueous sulfuric acid solu-
tions using different types of organic and inorganic
INTRODUCTION
Ceric sulfate has been widely employed as an oxi-
dant in the mechanistic studies in aqueous sulfuric acid
medium [1–8] because of its stability and its availabil-
Correspondence to: A. Fawzy; e-mail: afsaad13@yahoo.com.
ꢀC
2017 Wiley Periodicals, Inc.

144
FAWZY, ALTHAGAFI, AND ALTASS
substrates. In aqueous sulfuric acid solutions,
cerium(IV) can exist as a mixture of different types
of sulfate species such as Ce(SO4)2+, Ce(SO4)2,
HCe(SO4)3−, and H3Ce(SO4)4− [1–5].
per sulfate, silver nitrate, acrylonitrile, and methanol
used were of analytical grade with about 99.9%
purity.
Kinetic Measurements. The rates of disappearance
of cerium(IV) in aqueous sulfuric acid solution for
uncatalyzed and catalyzed reactions were followed
under conditions in which [His] ꢀ [Ce(IV)]. The
ionic strength was maintained constant (adjusted to
3.0 mol dm−3) by the addition of sodium sulfate elec-
trolyte. The oxidation reactions were carried out in a
thermostated cell compartment at a fixed temperature
(25°C), which was controlled to within 0.1°C. The
reactions were initiated by mixing thermally preequili-
brated solutions prepared under predetermined con-
ditions. Ce(IV) solution was then added to the re-
action vessel to initiate the oxidation process. The
progress of the reactions was followed by monitor-
ing the decay of cerium(IV) absorbance as a func-
tion of time at its absorption maximum, λ = 316 nm.
The absorbance measurements were conducted on a
temperature-controlled Shimadzu UV-VIS-NIR-3600
double-beam spectrophotometer (Japan). First-order
plots of ln(absorbance) versus time were straight
lines up to about 80% of the reaction completion.
The pseudo–first-order rate constants (kU and kC)
were determined as the slopes of these plots, and
average values of at least two independent kinetic
measurements were taken for calculation of the rate
constants.
The amino acid histidine (His) was used to treat
various diseases, and deficiency of it in hemoglobin
can cause poor hearing. Histidine is very important be-
cause the body uses it to manufacture of histamine,
which is responsible for a wide range of physio-
logical processes. It provides metal-binding sites on
many enzymes and is directly involved in catalysis.
It also forms complexes with some transition metal
ions in aqueous media [13,14]. The oxidation of his-
tidine by different oxidants has been previously stud-
ied in different media [13–19], and in most cases the
main oxidation product of histidine was 2-imidazole
acetaldehyde.
A literature survey revealed no work has been done
on the kinetics and mechanism of oxidation of histi-
dine by cerium(IV) in the absence or presence of a
catalyst. This leads us to investigate the kinetics and
mechanistic aspects of the oxidation of this amino acid
by cerium(IV) in aqueous sulfuric acid solutions in the
absence and presence of two different metal ions cata-
lysts, namely silver(I) and copper(II). We aim to inves-
tigate the selectivity of histidine toward cerium(IV) in
sulfuric acid solutions, to understand the active species
of the oxidant in this medium, to check the catalytic
activity of Ag(I) and Cu(II) catalysts toward histidine
oxidation, and to elucidate plausible reaction mecha-
nisms.
Whereas the reaction between cerium(IV) and his-
tidine in sulfuric acid solution proceeded with a mea-
surable rate in the absence of metal ions catalysts, the
catalyzed reactions were thought to proceed in a par-
allel path, with contributions from both uncatalyzed
and catalyzed oxidation reactions. Hence, the total rate
constant (kT) equals kU + kC.
EXPERIMENTAL
Materials and Methods
A few kinetic runs were carried out after bubbling
with purified nitrogen. The results of these runs were
compared with those conducted under air. The results
were the same, suggesting that the dissolved oxygen
did not affect the rate constants.
Materials. A stock solution of histidine was freshly
prepared by dissolving the required amount of the sam-
ple (Merck, Germany; 99.3%) in double-distilled wa-
ter. A solution of cerium(IV) was prepared by dissolv-
ing ceric ammonium sulfate (Sigma-Aldrich, Louis,
USA; ࣙ99%) in a 1.0 mol dm−3 sulfuric acid solu-
tion (Merck), then diluting with double-distilled water.
The concentration of cerium(IV) solution was deter-
mined by titrating it against ferrous ammonium sul-
fate solution. The solution of cerium(IV) was stored
in the dark and was used after 24 h, since hydroly-
sis is negligible or ruled out after 12 h of preparation
[20]. Cerium(III) solution was prepared by dissolv-
ing cerium(III) acetate (S.D. Fine Chem, Germany) in
double-distilled water. Other chemicals and reagents
such as sulfuric acid, sodium sulfate, hydrated cop-
RESULTS
Stoichiometry and Product Identification
The stoichiometry, determined spectrophotometrically
and by titration, indicates the consumption of two
Ce(IV) ions per one molecule of histidine yielding the
final oxidation products as illustrated in the following
equation:
International Journal of Chemical Kinetics DOI 10.1002/kin.21063

SILVER(I) AND COPPER(II) CATALYSIS FOR OXIDATION OF HISTIDINE BY CERIUM(IV)
145
H
N
H
N
N
N
+ CO2 + NH4+ + 2Ce3+ + 2SO42- + H+
+ 2Ce(SO4)2+ + H2O
(1)
OH
O
H
H2N
O
2-imidazole
Histidine
acetaldehyde
Equation (1) is consistent with the results of
To study the effect of the hydrogen ion concentra-
tion on the oxidation rates, kinetic runs were carried
out by varying the hydrogen ion concentration (0.5–
3.0 mol dm−3) using sulfuric acid and keeping other
reactants concentrations constant. The rates of both un-
catalyzed and catalyzed oxidation reactions were found
to increase with increasing [H+], as listed in Table I.
Plots of kU and kC versus [H+] were linear with positive
product analysis. The product, 2-imidazole acetalde-
hyde, was estimated quantitatively by its 2,4-
dinitrophenylhydrazine derivative [21]. Other products
were identified as discussed earlier [22]. Similar oxi-
dation products were obtained in the oxidation of his-
tidine by other oxidants [13–15,17].
Spectral Changes
1.0
Ce(IV)
The spectral changes during the oxidation of His
by cerium(IV) in sulfuric acid solutions in the pres-
ence of Ag(I) and Cu(II) catalysts are illustrated in
Figs. 1a and 1b. In both cases, the spectra indicate grad-
ual disappearance of the Ce(IV) band at its absorption
maximum as a result of its reduction to Ce(III). It can
be observed that the disappearance of Ce(IV) band in
the presence of Ag(I) occurs faster than that in the
presence of Cu(II).
0.8
(a)
0.6
0.4
Ag(I)
0.2
His
0.0
200
Order of Reactions
250
300
350
400
450
The order of both uncatalyzed and catalyzed reac-
tions with respect to the reactants was determined
from the slopes of the plots of logkU and logkC ver-
sus log(concentration) by varying the concentrations
of substrate, acid, and catalyst, in turn, while keeping
all other conditions constant.
The concentration of the oxidant, cerium(IV), was
varied in both reactions in the range of 0.5 × 10−4 to
5.0 × 10−4 mol dm−3, keeping all other variables con-
stant. The plots of ln(absorbance) versus time were
found to be linear to almost 80% of the reactions com-
pletion. Furthermore, an increase in the initial oxidant
concentration did not alter the oxidation rates of His
(Table I). These results indicate that the order of reac-
tions with respect to [Ce(IV)] is one.
The observed rate constants were determined at dif-
ferent initial concentrations of His, while other vari-
ables were kept constant. Increasing [His] increased
the oxidation rates (Table I). Plots of the observed rate
constants versus [His] were linear with positive inter-
cepts as shown in Fig. 2, confirming fractional-first-
order dependencies with respect to [His].
Wavelength, nm
1.0
Ce(IV)
0.8
0.6
0.4
0.2
0.0
(b)
Cu(II)
His
200
250
300
350
400
450
Wavelength, nm
Figure 1 Spectral changes during silver(I)- and copper(II)-
catalyzed oxidations of histidine (His) by cerium(IV) in sul-
furic acid solutions. [His] = 6.0 × 10−3, [Ce(IV)] = 2.0 ×
10−4, [H+] = 2.0, I = 3.0, [Ag(I)] = 8.0 × 10−6, and [Cu(II)]
= 8.0 × 10−4 mol dm−3 at 25°C. Scan time intervals = 1
min. [Color figure can be viewed at wileyonlinelibrary.com]
International Journal of Chemical Kinetics DOI 10.1002/kin.21063

146
FAWZY, ALTHAGAFI, AND ALTASS
Table I Effect of Varying [Ce(IV)], [His], [H+], [M(Y)], and Ionic Strength, I, on the First-Order Rate Constants
in the Uncatalyzed and Silver(I)- and Copper(II)-catalyzed Oxidations of Histidine by Cerium(IV)
in Sulfuric Acid Solutions at 25°C
105 k (s−1
)
Ag(I)
Cu(II)
104 [Ce(IV]
103 [His]
[H+]
*10X[M(Y)]
I
(mol dm−3
)
(mol dm−3
)
(mol dm−3
)
(mol dm−3
)
(mol dm−3
)
kU
kT
kC
kT
kC
0.5
1.0
2.0
3.0
4.0
5.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
6.0
6.0
6.0
6.0
6.0
6.0
2.0
4.0
6.0
8.0
10.0
12.0
6.0
6.0
6.0
6.0
6.0
6.0
6.0
6.0
6.0
6.0
6.0
6.0
6.0
6.0
6.0
6.0
6.0
6.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
0.5
1.0
1.5
2.0
2.5
3.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
2.0
8.0
8.0
8.0
8.0
8.0
8.0
8.0
8.0
8.0
8.0
8.0
8.0
8.0
8.0
8.0
8.0
8.0
8.0
2.0
4.0
6.0
8.0
10.0
12.0
8.0
8.0
8.0
8.0
8.0
8.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.3
3.6
3.9
4.2
4.5
62.3
64.4
64.0
65.7
63.9
63.4
33.8
49.7
64.0
76.1
85.0
94.9
23.0
37.4
51.1
64.0
81.0
90.9
64.0
64.0
64.0
64.0
64.0
64.0
64.0
67.2
70.8
74.7
78.3
81.4
279.9
288.9
290.3
293.8
294.3
294.0
165.1
235.2
290.3
339.7
410.3
449.8
119.3
173.8
233.0
290.3
352.2
414.9
118.3
175.0
244.9
290.2
345.6
428.1
290.2
301.9
316.9
328.9
340.3
352.7
217.6
224.5
226.3
228.1
230.4
230.6
131.3
185.5
226.3
263.6
325.3
354.9
96.3
136.4
181.9
226.3
271.2
324.0
54.3
111.0
180.9
226.2
281.6
364.1
226.2
234.7
246.1
254.2
262.3
271.3
182
119.7
124.0
122.4
125.2
121.6
120.3
71.2
188.4
186.4
190.9
185.5
183.7
105.0
149.1
186.4
220.4
255.3
301.1
73.2
112.3
139.4
186.4
222.2
264.2
96.3
122.1
154.5
186.4
220.2
242.7
186.4
194.3
202.1
211.1
217.7
225.4
99.4
122.4
144.3
170.3
206.2
50.2
74.9
88.3
122.4
141.2
173.3
32.3
58.1
90.5
122.4
156.2
178.7
122.4
127.1
131.3
136.2
139.4
144.0
Experimental error = 4%.
*X = 6.0 for Ag(I) and X = 4.0 for Cu(II).
slopes (Fig. 3), suggesting that the orders with respect
to [H+] were less than unity.
Effect of Ionic Strength and Dielectric
Constant
To evaluate the reactions orders with respect to the
concentrations of Ag(I) and Cu(II) catalysts, the ox-
idation rates were measured at various [M(Y)], (2.0–
12.0) × 10−6 mol dm−3 for Ag(I) and (2.0–12.0) ×
10−4 mol dm−3 for Cu(II), while other variables were
kept constant. The reaction rate increased as [M(Y)]
increased (Table I). The orders with respect to Ag(I)
and Cu(II) were less than unity, as determined from
the plots of log kC versus log [M(Y)], as shown in
Fig. 4.
The effect of ionic strength on the rates of uncatalyzed
and catalyzed reactions was studied by varying the
ionic strength in the range of 3.0–4.5 mol dm−3 using
sodium sulfate solution while maintaining the concen-
trations of all other reactants constant. Increasing ionic
strength of the medium increased the oxidation rates
of both reactions, and the Debye–Huckel plots were
linear with positive slopes as illustrated in Fig. 5a.
To determine the effect of dielectric constant (D)
of the medium on the oxidation rates, the reactions
International Journal of Chemical Kinetics DOI 10.1002/kin.21063

SILVER(I) AND COPPER(II) CATALYSIS FOR OXIDATION OF HISTIDINE BY CERIUM(IV)
147
400
300
200
100
0
-2.2
-2.4
Uncatalyzed reaction
Ag(I)-catalyzed reaction
Cu(II)-catalyzed reaction
-2.6
-2.8
-3.0
-3.2
-3.4
-3.6
Ag(I)
Cu(II)
-6.0
-5.5
-5.0
-4.5
-4.0
-3.5
-3.0
-2.5
0
2
4
6
8
10
12
log [M(Y)]
103 [His], mol dm-3
Figure 4 Plots of log kC versus log M(Y) in the metal ions–
Figure 2 Effect of [His] on the rate constants in the uncat-
alyzed and metal ions–catalyzed oxidations of histidine by
catalyzed oxidations of histidine by cerium(IV) in sulfuric
acid solutions. [His] = 6.0 × 10−3, [Ce(IV)] = 2.0 × 10−4
,
cerium(IV) in sulfuric acid solutions. [Ce(IV)] = 2.0 × 10−4
,
[H+] = 2.0, and I = 3.0 mol dm−3 at 25°C. [Color figure
[H+] = 2.0, I = 3.0, [Ag(I)] = 8.0 × 10−6 and [Cu(II)] =
8.0 × 10−4 mol dm−3 at 25°C. [Color figure can be viewed
at wileyonlinelibrary.com]
can be viewed at wileyonlinelibrary.com]
tent). Plots of log kU and log kC versus 1/D were linear
with negative slopes (Fig. 5b).
400
Uncatalyzed reaction
Ag(I)-catalyzed reaction
Cu(II)-catalyzed reaction
Effect of [HSO4−]
300
To study the effect of the bisulfate ion on the reaction
rate, the oxidation reaction was carried out at concen-
200
100
0
trations of bisulfate in the range of 1.0–3.5 mol dm−3
,
keeping the concentrations of all other reactants con-
stant. It was observed that increasing the bisulfate con-
centration decreases the reaction rate, and the order
with respect to [HSO4−] was negative. This is con-
firmed by the linear plot of 1/kU versus [HSO4−] with
a positive intercept, as shown in Fig. 6. Therefore,
[HSO4−] shows a rate-retarding effect.
0.0
0.5
1.0
1.5
2.0
2.5
3.0
[H+], mol dm-3
Figure 3 Effect of [H+] on the rate constants in the uncat-
alyzed and metal ions–catalyzed oxidations of histidine by
cerium(IV) in sulfuric acid solutions. [His] = 6.0 × 10−3
,
Effect of Initially Added Product
[Ce(IV)] = 2.0 × 10−4, I = 3.0, [Ag(I)] = 8.0 × 10−6, and
[Cu(II)] = 8.0 × 10−4 mol dm−3 at 25°C. [Color figure can
be viewed at wileyonlinelibrary.com]
The effect of added cerium(III) was studied over the
concentration range of 0.5 × 10−4 to 6.0 × 10−4 mol
dm−3 at fixed concentrations of other reactants. Ce(III)
did not have a significant effect on the oxidation rates
of both uncatalysed and catalyzed reactions.
were studied at different solvent compositions (vol%)
of acetic acid and water. At different compositions, the
dielectric constant of the medium was calculated using
the equation of Laidler and Eyring [23]: D = D1V1+
D2V2, where V1 and V2 are volume fractions, and D1
and D2 are dielectric constants of water and acetic
acid as 78.5 and 6.15 at 25°C, respectively. The rate
constants clearly decreased with decreasing dielectric
constant of the solvent (i.e., increasing acetic acid con-
Effect of Temperature
The oxidation rates of both uncatalyzed and metal
ions–catalyzed paths were measured at five different
temperatures (288–308 K) at constant concentrations
of the reactants. The results indicate that the rate con-
stants increased with raising temperature. The activa-
tion parameters of the second-order rate constants, kꢁ
International Journal of Chemical Kinetics DOI 10.1002/kin.21063

148
FAWZY, ALTHAGAFI, AND ALTASS
4
3
2
1
0
-5.5
-6.0
-6.5
-7.0
-7.5
Uncatalyzed reaction
Ag(I)-catalyzed reaction
Cu(II)-catalyzed reaction
(a)
0
1
2
3
4
[HSO4-], mol dm-3
0.63
0.64
0.65
0.66
0.67
0.68
0.69
I1/2 / (1+I1/2
)
Figure 6 A plot of 1/kU versus [HSO4−] in the uncatalyzed
oxidation of histidine by cerium(IV) in sulfuric acid solution.
[His] = 6.0 × 10−3, [Ce(IV)] = 2.0 × 10−4, [H+] = 2.0,
and I = 3.0 mol dm−3 at 25°C.
-2.4
-2.8
-3.2
-3.6
Uncatalyzed reaction
Ag(I)-catalyzed reaction
Cu(II)-catalyzed reaction
(b)
conditions, the tests were negative. These tests indi-
cated that the reactions proceeded through free radical
paths.
DISCUSSION
It was reported [13,14,24] that amino acids tend to
protonate at higher acid concentrations. In the present
investigation, the high concentration of hydrogen ions
used, as well as the significant increase in the oxida-
tion rates upon increasing the acid concentration sug-
gested protonation of histidine before its reaction with
the oxidant, as represented by Eq. (3) in Scheme 1.
Therefore, the protonated form of histidine (His+) is
suggested to be the kinetically reactive species in the
rate-determining step.
12
14
16
18
20
103 (1/
D)
Figure 5 Effect of (a) ionic strength and (b) dielectric con-
stant of the reaction medium on the rate constants in the
uncatalyzed and metal ions–catalyzed oxidations of histi-
dine by cerium(IV) in sulfuric acid solutions. [His] = 6.0 ×
10−3, [Ce(IV)] = 2.0 × 10−4, [H+] = 2.0, [Ag(I)] = 8.0 ×
10−6, and [Cu(II)] = 8.0 × 10−4 mol dm−3 at 25°C. [Color
figure can be viewed at wileyonlinelibrary.com]
On the other hand, in a sulfuric acid medium [3,4],
cerium(IV) forms several sulfate complexes having an
increasing number of coordinated sulfate ions accord-
ing to the equilibria shown in Eqs. (A)–(C) [20,25],
(kꢁ= k/[His], where k is kU or kC) were calculated using
Arrhenius (Fig. 7a) and Eyring (Fig. 7b) plots and are
listed in Table II.
Ce4+ + HSO4− ꢀ Ce (SO4)2+ + H+ Q1 = 3500
(A)
Test for Free Radical Intermediates
CeSO24+ + HSO4− ꢀ Ce (SO4)2 + H+ Q2 = 200
The generation of free radicals during a reaction was
examined by a polymerization test. Known quantities
of acrylonitrile scavenger were added to samples of
the reaction mixtures that had been kept in inert at-
mosphere for 2 h. White precipitates (as a result of
acrylonitrile polymerization) were formed upon dilut-
ing the reaction mixtures with methanol indicating free
radicals intervention in the reactions. When the experi-
ments were repeated in the absence of His under similar
(B)
Ce(SO4)2 + HSO−4 ꢀ Ce(SO4)32− + H+ Q3 = 20
(C)
The equilibrium constants, Q1, Q2, and Q3, are given
at an ionic strength of 2.0 mol dm−3 and 25°C.
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SILVER(I) AND COPPER(II) CATALYSIS FOR OXIDATION OF HISTIDINE BY CERIUM(IV)
149
0
-1
-2
-3
point to the fact that Ce(SO4)2+ may be the kinetically
active species for cerium(IV) sulfate complexes. This
case corresponds to Eq. (3) [20]. Also, because the rate
constants increased upon increasing the ionic strength
and dielectric constant of the reactions media, it is
likely that Ce(SO4)2+ is one of the reactive species in
these reactions, as reactions were suggested [26,27] to
occur between two similarly charged species, that is,
(a)
Uncatalyzed reaction
Ag(I)-catalyzed reaction
Cu(II)-catalyzed reaction
between His+ and Ce(SO4)2+
.
In earlier work [28], it was reported that a reaction
rate that decreases as the concentration of HSO4− in-
creases means that Ce(SO4)2+ and Ce(SO4)2 are kinet-
ically active in the reaction, but the kinetic contribution
of Ce(SO4)2 is negligibly small and the major active
species is Ce(SO4)2+. Furthermore, electron flow to-
ward the Ce(IV) center in Ce(SO4)2+ is favored over
Ce(SO4)2 because the Ce(IV) centered in Ce(SO4)2+
bears a more positive charge. Ce(SO4)2+ species has
also been suggested in earlier works [28–31].
3.25
3.30
3.35
3.40
3.45
3.50
103 (1/T), K-1
-29
-30
-31
-32
-33
(b)
Uncatalyzed reaction
Ag(I)-catalyzed reaction
Cu(II)-catalyzed reaction
Mechanism of Uncatalyzed Oxidation
The results of the kinetic study of the oxidation of
His by cerium(IV) in sulfuric acid medium suggests
formation of an intermediary complex between the
protonated histidine, His+, and the kinetically active
species of Ce(IV), Ce(SO4)2+. Formation of an inter-
mediary complex is supported by the less than unit or-
der with respect to [His]. Complex formation was also
proved kinetically by the nonzero intercept of the plot
of 1/kU versus 1/[His] (Fig. 8). This kinetic result fa-
vors the possible formation of an intermediary complex
between the oxidant and substrate [32]. Furthermore,
complex formation was supported by the obtained neg-
ative value of the entropy of activation, ࢞S# = −94.04
J K−1 mol−1, which corresponds to the decrease in
the degrees of freedom accompanying the transition
from the initial state to the transition state [33,34]. The
3.25
3.30
3.35
3.40
3.45
3.50
103 (1/T) K-1
Figure 7 (a) Arrhenius and (b) Eyring plots in the uncat-
alyzed and metal ions–catalyzed oxidations of histidine by
cerium(IV) in sulfuric acid solutions. [His] = 6.0 × 10−3
,
[Ce(IV)] = 2.0 × 10−4, [H+] = 2.0, I = 3.0, [Ag(I)] = 8.0
× 10−6, and [Cu(II)] = 8.0 × 10−4 mol dm−3. [Color figure
can be viewed at wileyonlinelibrary.com]
Based on the effect of [H+] and [HSO4−] on the
rate of the oxidation, one of these species was treated
as the kinetically active species for cerium(IV). The
observation in our work that HSO4− ions retard the rate
of oxidation while increasingthe rate as [H+] increases,
positive values of both ꢀH and ꢀG indicate en-
dothermic formation of the intermediary complex and
its nonspontaneity, respectively.
ࣔ
ࣔ
The oxidation reaction is proposed to be a one-
electron process in the rate-determining step, which
Table II Activation Parameters of kꢁ in the Uncatalyzed and Metal Ions–Catalyzed Oxidations of Histidine by
Cerium(IV) in Sulfuric Acid Solutions. [His] = 6.0 × 10−3, [Ce(IV)] = 2.0 × 10−4, [H+] = 2.0, I = 3.0, [Ag(I)] = 8.0 ×
10−6, and [Cu(II)] = 8.0 × 10−4 mol dm−3
ࣔ
ࣔ
ࣔ
ࣔ
Reaction
ꢀS (J mol−1K−1
)
ꢀH (kJ mol−1
)
ꢀG
(kJ mol−1
)
Ea (kJ mol−1
)
298
Uncatalyzed reaction
Ag(I)-catalyzed reaction
Cu(II)-catalyzed reaction
−94.04
−54.42
−76.71
60.46
27.21
39.30
88.47
43.43
62.16
69.09
36.36
47.27
Experimental error 3%.
International Journal of Chemical Kinetics DOI 10.1002/kin.21063

150
FAWZY, ALTHAGAFI, AND ALTASS
K
2+
-
CeSO4 + HSO4
Ce(SO4)2 + H+
(2)
(3)
H
N
H
N
H
N
N
N
N
H+
-
OH
O
OH
+
+
K1
H3N
H2N
H3N
O
O
O
H
N
H
N
N
N
H
K2
+ Ce(SO4)2+
Ce(SO4)2+
OH
+
O
+
(4)
(5)
(6)
H3N
H3N
O
O
(C1)
H
N
H
N
N
N
H
k1
slow
.
Ce(SO4)2+
+ Ce3+ + SO42- + H+
O
+
O
+
H3N
H3N
O
O
H
N
H
N
N
N
fast
.
+ CO2 + H+
O
+
.
H2N
H3N
H
O
H
N
H
N
N
N
fast
+ Ce3+ + SO42-
+ Ce(SO4)2+
+
(7)
(8)
.
H2N
H
H2N
H
H
N
H
N
N
N
fast
+ NH4+
+ H2O
+
H2N
O
H
H
Scheme 1 Mechanism of uncatalyzed oxidation of histidine by cerium(IV) in sulfuric acid solution.
International Journal of Chemical Kinetics DOI 10.1002/kin.21063

SILVER(I) AND COPPER(II) CATALYSIS FOR OXIDATION OF HISTIDINE BY CERIUM(IV)
151
1/[H+], dm3 mol-1
Comparing Eqs. (9) and (10), we obtain
k1K1K2[His][H+]
1 + K1[H+] + K1K2[His][H+] + K[HSO4−](K1 + [H+]−1
1.2
2.0
0.0
0.4
0.8
1.6
5000
4000
3000
2000
1000
0
kU
=
)
(11)
and with rearrangement of Eq. (11), Eqs. (12) and (13)
are obtained:
ꢀ
ꢁ
1
kU
1 + K1[H+]
1
1
=
+
+ KꢁK[HSO−4 ]
k1K1K2[H+] [His] k1
(12)
ꢀ
ꢁ
0
100
200
300
400
500
1
kU
1
1
1
=
+
1/[His], dm3 mol-1
k1K1K2[His] [H+] k1K2[His]
1
Figure 8 Plots of 1/kU versus 1/[His] and 1/[H+] for the
uncatalyzed oxidation of histidine by cerium(IV) in sulfuric
acid solution. [Ce(IV)] = 2.0 × 10−4 and I = 3.0 mol dm−3 at
25°C. [Color figure can be viewed at wileyonlinelibrary.com]
+
+ KꢁK[HSO−4 ]
(13)
k1
K1 + [H+]
where Kꢁ =
k1K1K2[His][H+]
Regarding to Eqs. (12) and (13), the plots of 1/kU versus
1/[His] at constant [H+] and 1/kU versus 1/[H+] at
constant [His] should be linear with positive intercepts
and are found to be so as shown in Fig. 8.
leads to decomposition of the intermediary complex
forming an intermediary radical and Ce(III) ion. The
formation of a free radical is evident through the acry-
lonitrile polymerization test. The rate was not affected
by Ce(III), suggesting that the reaction step involv-
ing decomposition of the intermediary complex into
a histidine free radical and Ce(III) ion may not be a
reversible process. The rate-determining step should
be irreversible, which is generally the case for one-
electron oxidants [35]. Next, decarboxylation of the
histidine free radical occurred, forming a new radical
intermediary (X•). The latter was rapidly attacked by
another Ce(SO4)2+ species to yield the final oxidation
Mechanism of Metal Ions–Catalyzed
Oxidations
The kinetics of the oxidations of His by cerium(IV)
in sulfuric acid solutions in the presence of small
amounts of silver(I) and copper(II) catalysts are sim-
ilar to that for the uncatalyzed oxidation reaction. In
addition, these reactions show first order with respect
to silver(I) and copper(II) concentrations. Less than
unit orders for the concentrations of His and metal ion
catalysts suggest formation of complexes between His
and metal ions catalysts [3,12–14,36] prior to reac-
tion with the oxidant. Complexes formation was also
proved kinetically by the nonzero intercepts of the plots
of [M(Y)]/kC versus 1/[His] (M = metal; Y = oxida-
tion state of M) (Fig. 9). Such complexes have earlier
been reported between His and silver(I) [13] and be-
tween His and copper(II) [14]. These complexes react
in slow steps with Ce(SO4)2+ to give rise to the in-
termediary radical and Ce(III) ion with regeneration of
the catalyst. This is followed by decarboxylation of His
free radical, forming a new radical intermediary (X•).
The latter is rapidly attacked by another Ce(SO4)2+
species to yield the final oxidation products as illus-
trated in Scheme 2.
products. The suggested mechanism is illustrated in
Scheme 1.
The suggested mechanism leads to the following
rate-law expression (see the Appendix),
k1K1K2[Ce(IV)][His][H+]
Rate =
1 + K1[H+] + K1K2[His][H+] + K[HSO4−](K1 + [H+]−1
)
(9)
The rate law (9) is consistent with all the observed
orders with respect to different species.
Under pseudo–first-order conditions, the rate law
can be expressed by Eq. (10),
−d[Ce(IV)]
Rate =
= kU [Ce (IV)]
(10)
dt
International Journal of Chemical Kinetics DOI 10.1002/kin.21063

152
FAWZY, ALTHAGAFI, AND ALTASS
H
N
H
N
N
N
H
O
K3
OH + M(Y)
+
+
M(Y)
C2
(14)
H3N
H3N
O
O
(
)
H
N
H
N
N
N
+ M(Y) + Ce3+
+ SO42- + H+
H
k2
slow
.
+ Ce(SO4)2+
O
O
+
M(Y)
+
(15)
(16)
H3N
H3N
O
O
H
N
H
N
N
N
fast
.
+ CO2 + H+
O
+
.
H2N
H3N
H
O
H
N
H
N
N
N
fast
+ Ce3+ + SO42-
+ Ce(SO4)2+
(17)
(18)
+
.
H2N
H2N
H
H
H
N
H
N
N
N
fast
+ NH4+
+ H2O
+
H2N
O
H
H
Scheme 2 Mechanism of metal ions-catalyzed oxidations of histidine by cerium(IV) in sulfuric acid solutions.
An alternative oxidation mechanism [3,12,36] for
ion. Such complex rapidly decomposes giving rise
to the intermediary radical with regeneration of the
catalyst M(Y), followed by subsequent fast steps to
yield the final oxidation products as illustrated in
metal ions–catalyzed oxidations may be proposed. It
involves formation of an intermediary complex be-
tween the metal ion catalyst, M(Y), and the amino
acid (C2), that on further interaction with Ce(IV) in
the rate-determining step yields another complex (C3)
of a higher valence metal ion, M(Y+1), and Ce(III)
Scheme 3.
In a similar manner to the uncatalyzed reac-
tion, the proposed mechanism leads to the following
International Journal of Chemical Kinetics DOI 10.1002/kin.21063

SILVER(I) AND COPPER(II) CATALYSIS FOR OXIDATION OF HISTIDINE BY CERIUM(IV)
153
H
N
H
N
N
N
H
K3
(19)
OH
+ M(Y)
O
M(Y)
C2
+
+
H3N
H3N
O
O
(
)
H
N
H
N
N
N
H
O
H
O
k3
slow
+
Ce(SO4)2+
+
Ce3+ SO42-
+
(20)
+
M(Y)
+
M(Y+1)
H3N
H3N
O
O
(C2)
(C3)
H
N
H
N
N
N
H
O
fast
.
(21)
(22)
+
M(Y) + H+
O
+
M(Y+1)
+
H3N
H3N
O
O
H
N
H
N
N
N
N
N
fast
.
+ CO2 + H+
O
+
.
H2N
H3N
H
O
H
N
H
N
N
fast
(23)
(24)
+ Ce3+ + SO42-
+ Ce(SO4)2+
+
.
H2N
H2N
H
H
H
H
N
N
N
fast
+ NH4+
+ H2O
+
H2N
O
H
H
Scheme 3 An alternative mechanism of metal ions–catalyzed oxidations of histidine by cerium(IV) in sulfuric acid solutions.
rate law expression shown in the following
equation:
The rate law expressed in Eq. (25) is consistent with all
the observed orders with respect to different species.
Under pseudo–first-order conditions, the rate law can
be expressed by Eq. (26),
k2K1K3[Ce(IV)][His][H+][M(Y)]
Rate =
1 + K1[H+] + K1K3[His][H+] + K[HSO4−](K1 + [H+]−1
)
−d[Ce(IV)]
Rate =
= kC [Ce (IV)]
(26)
(25)
dt
International Journal of Chemical Kinetics DOI 10.1002/kin.21063

154
FAWZY, ALTHAGAFI, AND ALTASS
ꢀ
ꢁ
1.4
1.2
1.0
0.8
0.6
0.4
0.2
0.0
1.4
1.2
1.0
0.8
0.6
0.4
0.2
0.0
[M(Y)]
kC
1
1
1
=
+
k2K1K3[His] [H+] k2K3[His]
(R 2 = 0.989)
1
+
+ KꢁꢁK[HSO−4 ]
(29)
k2
K1 + [H+]−1
where Kꢁꢁ =
k2K1K3[His][H+]
(R 2 = 0.993)
According to Eqs. (28) and (29), plots of [M(Y)]/kC
versus 1/[His] at constant [H+], and [M(Y)]/kC against
1/[H+] at constant [His], should be linear with positive
intercepts on the [M(Y)]/kC axes. The experimental
results satisfy these requirements, as shown in Figs. 9
and 10, respectively.
The mechanism of catalysis by metal ions can be
quite complicated due to formation of different inter-
mediate complexes. Although the mechanism of catal-
ysis depends on the nature of the substrate, oxidant, and
the experimental conditions, it has been shown [14] that
metal ions act as catalysts by a few different paths such
as the formation of complexes with reactant, or direct
oxidation of the substrate, or through the formation of
free radicals.
0
100
200
300
400
500
1/[His], dm3 mol-1
Figure 9 Plots of [Ag(I)]/kC and [Cu(II)]/kC versus 1/[His]
in the metal ions–catalyzed oxidations of histidine (His) by
cerium(IV) in sulfuric acid solutions. [Ce(IV)] = 2.0 × 10−4
,
[H+] = 2.0, and I = 3.0 mol dm−3 at 25°C. [Color figure
can be viewed at wileyonlinelibrary.com]
2.0
1.5
1.0
0.5
0.0
2.0
1.5
1.0
0.5
0.0
(R 2 = 0.987)
CONCLUSIONS
A comparative study of the oxidations of histidine by
cerium(IV) in aqueous sulfuric acid solutions in the ab-
sence and presence of silver(I) and copper(II) catalysts
was performed. The oxidation reaction was found to
occur with a slow rate in the absence of the catalysts.
The catalytic efficiency of Ag(I) was higher than that
of Cu(II). In both cases, the final oxidation products of
histidine were identified as 2-imidazole acetaldehyde,
ammonium ion, and carbon dioxide.
(R 2 = 0.991)
0.0
0.5
1.0
1.5
2.0
1/[H+], mol dm
-3
Figure 10 Plots of [Ag(I)]/kC and [Cu(II)]/kC versus
1/[H+] in the metal ions–catalyzed oxidations of histidine
(His) by cerium(IV) in sulfuric acid solutions. [His] = 6.0 ×
10−3, [Ce(IV)] = 2.0 × 10−4, and I = 3.0 mol dm−3 at 25°C.
[Color figure can be viewed at wileyonlinelibrary.com]
APPENDIX: DERIVATION OF RATE-LAW
EXPRESSION OF THE UNCATALYZED
OXIDATION REACTION
Comparing Eqs. (25) and (26), we obtain Eq. (27),
k2K1K3[His][H+][M(Y)]
kC
=
1 + K1[H+] + K1K3[His][H+] + K[HSO4−](K1 + [H+]−1
)
According to the proposed mechanistic Scheme 1,
−d[Ce(IV)]
(27)
Rate =
= k1 [C1]
(A1)
dt
With rearrangement of Eq. (27), the following two
equations are obtained:
According to reactions (2)–(4),
[Ce(SO4)2][H+]
K =
ꢂ
ꢃ
ꢀ
ꢁ
, Ce (SO4)2+
[M(Y)]
kC
1 + K1[H+]
1
[Ce(SO4)2+][HSO4−]
=
k2K1K3[H+] [His]
[Ce(SO4)2][H+]
1
(A2)
+
+ KꢁK[HSO−4 ]
(28)
=
K[HSO−4 ]
k2
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SILVER(I) AND COPPER(II) CATALYSIS FOR OXIDATION OF HISTIDINE BY CERIUM(IV)
155
Under pseudo–first-order conditions,
[His+]
[His][H+]
ꢂ
ꢃ
ꢂ
ꢃ
K1 =
, His+ = K1 [His] H+
−d[Ce(IV)]
Rate =
= kU [Ce (IV)] (A13)
dt
(A3)
(A4)
Comparing Eqs. (A12) and (A13), we obtain Eq. (A14),
[C1]
K2 =
, [C1]
[Ce(SO4)2+][His+]
k1K1K2[His][H+]
1 + K1[H+] + K1K2[His][H+] + K[HSO4−](K1 + [H+]−1
ꢂ
ꢃ
KU
=
= K2 Ce (SO4)2+ [His+]
)
(A14)
K1K2[Ce(SO4)2][His][H+]2
ꢀ
ꢁ
=
K[HSO−4 ]
1
kU
1 + K1[H+]
1
1
=
+
+ KꢁK[HSO−4 ]
k1K1K2[H+] [His] k1
Substituting Eq. (A4) into Eq. (A1) leads to Eq. (A5).
(A15)
k1K1K2[Ce(SO4)2][His][H+]2
ꢀ
ꢁ
Rate =
(A5)
1
kU
1
1
1
K[HSO−4 ]
=
+
k1K1K2[His] [H+] k1K2[His]
1
The total concentration of histidine is given as follows:
+
+ KꢁK[HSO−4 ]
(A16)
k1
ꢂ
ꢃ
[His]T = [His]F + His+ + [C1]
(A6)
BIBLIOGRAPHY
[His]T
1 + K1[H+] +
[His]F =
(A7)
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]
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[His]F =
(A8)
1 + K1[H+]
Also,
ꢂ
ꢃ
[Ce (IV)]T = [Ce (SO4)2+] + Ce (SO4)2 + [C1] (A9)
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[H+] + K[HSO4−] + K1K2[His][H+]2
ꢂ
ꢃ
Ce (SO4)2
=
F
(A10)
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(A11)
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)
(A12)
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International Journal of Chemical Kinetics DOI 10.1002/kin.21063