The kinetics of cerium(IV) sulfate reaction with citrate and the thermodynamic characteristics of formation of intermediate complexes

Article abstract of DOI:10.1134/S0036024409060132

Intermediate cerium(IV)-citrate complexes formed at the first stage of the oxidation of citric acid (Citr) with cerium(IV) were studied spectrophotometrically and pH-potentiometrically at ionic strength I = 2 (sulfate medium). Their composition and the fo

Full text of DOI:10.1134/S0036024409060132

ISSN 0036-0244, Russian Journal of Physical Chemistry A, 2009, Vol. 83, No. 6, pp. 945–950. © Pleiades Publishing, Ltd., 2009.
Original Russian Text © O.O. Voskresenskaya, N.A. Skorik, 2009, published in Zhurnal Fizicheskoi Khimii, 2009, Vol. 83, No. 6, pp. 1079–1084.
PHYSICAL CHEMISTRY
OF SOLUTIONS
The Kinetics of Cerium(IV) Sulfate Reaction with Citrate
and the Thermodynamic Characteristics of Formation
of Intermediate Complexes
a
b
O. O. Voskresenskaya and N. A. Skorik
a
Joint Institute for Nuclear Research, Dubna, Moscow oblast, 141980 Russia
b
Tomsk State University, Tomsk, 634050 Russia
e-mail: voskr@jinr.ru
Received April 29, 2008
Abstract—Intermediate cerium(IV)–citrate complexes formed at the first stage of the oxidation of citric acid
(
Citr) with cerium(IV) were studied spectrophotometrically and pH-potentiometrically at ionic strength I = 2
(
sulfate medium). Their composition and the form of the organic ligand present in them, the thermodynamic
parameters of their formation, and the kinetic parameters of intramolecular redox decomposition were deter-
2
–
4
+
mined. A detailed scheme of the initial stages of the redox process in the Ce –SO –Citr system was consid-
4
ered, and the law of its initial rate and intermediate mechanism were determined. The results were compared
with the corresponding data on several oxycarboxylic acids and polyhydric alcohols. The inverse linear corre-
lation was found between the logarithms of stability constants and the logarithms of rate constants for intramo-
+
lecular redox decomposition of [(CeOH)H R] complexes with dibasic ligands of the type R = H L, H(OH)L,
–
2
2
and L(OH) . The stabilizing role played by ligand oxy groups in these complexes was demonstrated.
2
DOI: 10.1134/S0036024409060132
INTRODUCTION
The detailed mechanism of the BZh reaction is very
complex and includes dozens of intermediate com-
pounds. Its study is based on the so-called Field–
Koros–Noyes (FKN) theory [32] and its subsequent
developments [33–35]; it was the object of numerous
studies [1–4, 7, 9, 31–36]. The radical mechanism [2, 3]
and the intermediate and final products of the oxida-
Interactions of cerium(IV) with organic compounds
are extensively studied [1–19]. These studies are of
interest from both kinetic [1–22] and thermodynamic
[
5, 8, 23–26] points of view and are of considerable
importance in relation to the problem of the stabiliza-
tion of unstable oxidation states of metals [26–28] and
numerous applications of cerium(IV) in various
domains of chemistry as a complex-forming agent on
the one hand and a one-electron oxidizer on the other
tion of malonic acid CH (COOH) [3, 4] and its
2
2
derivatives [1, 4], oxalic (COOH) [3, 7, 38], citric
2
HOOCCH C(OH)(COOH)CH COOH [18–21], malic
2
2
HOOCHC(OH)CH COOH, tartaric (CH(OH)COOH)
[1, 25].
2
2
[
3, 5] and other carboxylic acids by cerium (IV) as the
The oxidation of aliphatic organic substances with
first stage of BZh reaction (1) have been actively stud-
ied. The possibility of formation of intermediate orga-
nocerium(IV) complexes in the BZh reaction was men-
tioned in several works [1, 4, 31, 32]. However, most of
these complexes have not been identified and their
structure, reactivity, and thermodynamic stability have
not been determined.
an active methylene group by cerium (the first stage of
the Belousov–Zhabotinskii (BZh) autooscillating reac-
tion) [1–4, 7, 9, 29–37] is of special interest. Concen-
tration oscillations in the BZh reaction occur because of
the alternation of two stages schematically written as
Ce(IV) + R = Ce(III) + P,
(1)
–
Ce(III) + BrO = Ce(IV) + P',
(2)
3
It is known from kinetic studies that polar organic
compounds (oxycarboxylic acids, alcohols, ketones,
where R is the reducing agent and P and P' are the reaction
products [31]. Oxycarboxylic and dicarboxylic acid (citric etc.) are oxidized by variable-valence metal ions
[
3, 29–31], malic [3, 30], oxalic [3, 7], malonic [2, 3], and (Ce(IV), Mn(III), …) according to the one-electron
methylmalonic [1, 4]), β-diketonates [28], saccharides [4], mechanism, which as a rule includes two stages. First,
alcohols [6], etc. can play the role of reducing agents. Oxi- an organic compound enters into the coordination
dizers can be cerium(IV) and one-electron oxidizers (e.g., sphere of the ion-oxidizer in an equilibrium reaction,
Mn(III)) close to it in the redox potential value and the and a redox process then occurs in this complex, which
kinetics of redox reactions [30].
is the rate-limiting stage [1, 4, 20–22, 31, 32],
9
45

9
46
VOSKRESENSKAYA, SKORIK
.
ric titration with a solution of NaOH which did not con-
0
2
0
3
–1
∆
D × 10
–D × 10 , s
2
–
tain the CO ion. The content of cerium(IV) in the salt
3
and freshly prepared solution was determined by back
titration with Mohr’s salt in the presence of ferroin [39]
before and after measurements. The optical density of
solutions was recorded on a SPECORD UV VIS spec-
trophotometer in the extinction measurement mode at a
constant wave number (λ = 400 nm) and a KF-5 photo-
electrocolorimeter with a MEA-4 recording device.
The reaction time was counted from the moment of
turning a vessel-mixer, into which the initial compo-
nents were placed.
4
4
3
2
1
0
0
0
0
40
30
20
10
2
0
∆D
.
0
–
4
D
3
1
1
2
3
c /c
The initial optical density of the reaction mixture of
R
M
0
the metal and ligand D was found by the linear extrap-
0
.2 0.4 0.6 0.8 1 – y
olation of kinetic curves in the logD –τ semilogarith-
0
0
˙
mic coordinates to the τ = 0 time moment, because
changes in the logarithm of the optical density in time
obeyed a straight line equation. The initial rate of the
0
Fig. 1. (1, 2) ∆D —composition and (3, 4) –D —compo-
sition diagrams: (1, 3) isomolar series of solutions (c =
.00 × 10 mol/l, pH 1.44, I = 2, T = 23.0°C, λ = 400 nm),
y is the mole fraction of the ligand and (2, 4) series of molar
ratios (cM = 2.8 × 10 mol/l, pH 2.23, I = 2, T = 25.0°C,
λ = 400 nm).
−
2
1
˙
–1
–
3
observed redox process –D (s ) was estimated graph-
ically in the same coordinates from the slope tanα =
0
i
0
i
(
D – D )/(τ – τ ) = const and calculated by the method
of least squares. Solutions were held at a constant tem-
perature using a Wobser thermostat. pH of mixtures of
components and solutions of cerium(IV) sulfate was
measured by a DATA METER precision pH meter.
Calculations were performed using the logarithms of
n + 1
K'
n + 1
k
n+
+
M
+ HR'
HR'…M
M + ·R' + H ,
(3)
K' = k /k , k ꢀ k > k.
1
–1
1
–1
However, kinetic data cannot be used to draw con-
clusions about the structure of this complex; we can
only estimate the effective constants of its formation at
a fixed pH value. These constants correspond to indefi-
nite forms of the metal and ligand in the complex. Ther-
modynamic methods of study of preequilibrium (3)
i
protonization constants B = [R]/([H R][H] ) of the
i
–i
citric acid (H R) anion, logB = 4.17 and logB =
–
2
1
2
1
7
.02 [40].
[
23–25] cannot be used to determine the kinetic param-
RESULTS AND DISCUSSION
The composition of intermediate complexes formed
eters of these complexes. Only kinetic generalizations
of thermodynamic methods [5, 8], on the one hand, pro-
vide the possibility of directly determining the compo-
2
4
–
4
+
in the Ce –SO –Citr system at the moment of mixing
sition of intermediate complexes, the form in which reaction mixture components was determined from the
organic ligands are present in them, and noneffective property–composition diagrams by the generalized
constants of their formation, which allows us to refine molar ratio and isomolar series methods [5, 8]. Accord-
scheme (3). On the other hand, they can be used to ing to Fig. 1 (lines 1, 2), the position of an extremum on
study the reactivity of these complexes and the kinetics the axis of compositions corresponds to the formation
and mechanism of the initial stages of the redox process in the system of a 1 : 1 cerium(IV)–citrate complex at
that occurs in the system.
In this work, this approach is applied to study the
cerium(IV)–citrate reaction in a sulfuric acid medium.
the initial time. The initial rate of the redox process,
2
which is a function of the yield of the complex, was
1
Here and throughout, the index –i(–x) in H R(H R) formulas
2
4
–
–i
–x
4
+
The results of earlier studies of the Ce –SO –Citr
denotes the number of dissociated protons (protons displaced
from reducing molecule R by complex formation).
system were reported in [22–24]. In the context of more
complex systems, this subsystem was considered in [3,
2
Under the conditions of the predominance of one intermediate
complex, the initial rate of the observed redox process is deter-
1
7–19, 29–31].
0
0
0
˙
mined by the equation –D = kobs
c
= kobsα c at cM < cR
C M
C
0
0
C
0
C
EXPERIMENTAL
We used doubly recrystallized citric acid of kh. ch.
chemically pure) grade and cerium(IV) sulfate
Ce(SO ) · 4H O of ch. d. a. (pure for analysis) grade.
The I = 2 ionic strength of solutions was adjusted using
ammonium sulfate of ch. d. a. grade. The concentration
of solutions of the organic acid was refined by pH met-
˙
and –D = kobs
c
= kobsα c at c < c , where c and c are
R
R
M
M
R
0
M
the initial concentrations of the metal ion and ligand; α
=
(
0
∞
0
C
0
∆D /(D – DM), c is the yield and concentration of the complex
4
2
2
0
∞
at time τ = 0; D , and D are light absorption by the reaction mix-
M
0
C
0
0
ture at α = 1 and metal, respectively; and ∆D = D – D [5].
M
RUSSIAN JOURNAL OF PHYSICAL CHEMISTRY A Vol. 83 No. 6 2009

THE KINETICS OF CERIUM(IV) SULFATE REACTION
947
determined by the generalized isomolar series and
.
0
3
–1
–
5
D × 10 , s
logβef
molar ratio methods. This rate reached a maximum
simultaneously with the maximum yield of the complex
at the stoichiometric ratio between the reagents. This
was also evidence of the formation of a 1 : 1 intermedi-
ate complex in the system (Fig. 1, lines 3, 4). Note that
0
x = 3
5
0
4
3
2
1
4
3 D , D
4
3
2
1
0
M
˙
0
the position of an extremum in the –D , D —composi-
tion diagrams remains constant until the depth of the
observed redox process reaches ξ ~ 0.9. The data on the
composition of the complex are in agreement with the
results obtained in [24] and can be considered a proof
of the validity of the suggestion of the formation of a
0.6
0.4
0.2
0
5
2
1
: 1 intermediate complex in the oxidation of citric
x = 2
1
acid by cerium(IV) [22, 31].
kobs
The form in which the coordinated ligand was present
in the complex and the stability constant of the complex
were determined by analyzing the property–medium
pH diagram (Fig. 2, lines 1–4) and by additionally ana-
lyzing the kinetic data. The predominant form of
cerium(IV) against the sulfate background at pH stud-
0
1
2
3
4
4
pH
0
1
2
3
c0 × 10 , mol/l
3
C
0
Fig. 2. (1) DM – pH, (2) D – pH, and (3) logβ – pH dia-
ef
n + 1
3+
ied is the monohydroxo form M
= CeOH [41–43].
2–
4
grams for the Ce –SO –Citr system (cM = 3.0 × 10−3 mol/l,
cR = 3.0 × 10–3 mol/l, I = 2, T = 25.0°C, λ = 400 nm),
4+
The number of protons x displaced from the R = Citr
molecule by the cerium(IV) ion when equilibrium
2
4
–
(
4) logβef – pH diagram for the Ce4+–SO –Malt system
(cM = 2.1 × 10–3 mol/l, cR = 2.0 × 10–3 mol/l, I = 2, T =
3
+
K
3 – x
+
CeOH + Citr
[(CeOH)H Citr]
+ xH ,
–
x
0
0
C
x
(4)
˙
1
7.7°C, λ = 400 nm), and (5) D – c diagram for the
[
(CeOH)H Citr][H]
–
x
x
K = ------------------------------------------------------ = β [H]
ef
2–
Ce4+–SO4 –Citr system (cM = 2.8 × 10–3 mol/l, pH 2.23,
I = 2, T = 25.0°C, λ = 400 nm).
[
CeOH][Citr]
was established (charges are omitted for simplicity)
was estimated graphically as the slope of the depen-
dence of logβ on pH,
The corresponding equilibrium constant
ef
0
C
[
(CeOH)H Citr]
c
–
2
logβef = logK + xpH.
(5)
β = ------------------------------------------ = ------------------------------------------- f
2
0 0
[
CeOH][H Citr]
–2
(c – c )(c – c )
M C R C
It follows from an analysis of the dependence of logβef
= β f ,
ef
2
(7)
3
+
on pH (Fig. 2, line 3) that the CeOH ion displaces two
protons (x = 2) from the citric acid molecule as a result
of complex formation. Most probably, two carboxylic
groups of citric acid participate in complex formation.
The result obtained for the form in which citric acid is
present in the cerium(IV)–citrate complex proves the
2
i
f 2 = 1 +
B [H] ,
i
∑
i = 1
was calculated for every point of the logβ = logβef
+
log f dependence. The logβ values were then aver-
2
0
validity of the suggestion made in [22] of the chelate aged over D – pH data series and the ascending portion
structure of the intermediate complex in the of the molar ratio diagrams (table). The logβ value
cerium(IV)–citrate reaction.
found for the cerium(IV)–citrate complex against the
sulfate background by the D – pH and molar ratio
0
According to the data obtained, complex formation
methods was logβ = 5.34 ± 0.15 (table). The K =
m
2
4
–
4
+
preequilibrium in the Ce –SO –Citr system can be
–1
0
.67 ≈ βef value estimated in [30] is an effective value
represented as
corresponding to an indefinite form of the presence of
the reagents in the cerium(IV)–citrate complex. The
3
+
2–
β
+
3
– x
CeOH + H Citr
[(CeOH)H Citr] . (6) logβ value obtained for the [(CeOH)H R]
com-
–
2
–2
–x
RUSSIAN JOURNAL OF PHYSICAL CHEMISTRY A Vol. 83 No. 6 2009

9
48
VOSKRESENSKAYA, SKORIK
Stability constants and rate constant for the redox decomposition of the cerium(IV)–citrate complex
D0
DM
pH
logβef
logβ
D0
–log(–D )
0
c /c
logβ
–logkC
˙
R
M
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
.227
.235
.255
.292
.350
.370
.415
.420
.430
.440
.460
.465
.495
.515
.575
.583
.585
0.207
0.208
0.210
0.208
0.208
0.191
0.177
0.174
0.161
0.161
0.142
0.132
0.111
0.102
0.069
0.074
0.069
1.52
1.61
1.73
1.90
2.06
2.21
2.28
2.33
2.37
2.40
2.45
2.53
2.58
2.66
2.85
2.90
2.91
1.17
1.31
1.57
1.86
2.28
2.45
2.71
2.74
2.82
2.87
3.01
3.06
3.25
3.38
3.83
3.89
3.92
5.17
5.14
5.18
5.13
5.24
5.14
5.27
5.22
5.23
5.22
5.28
5.19
5.30
5.30
5.46
5.45
5.46
0.285
0.330
0.375
0.410
0.450
0.475
0.500
0.515
0.519
0.525
0.525
0.525
0.539
0.530
0.550
0.545
0.550
2.89
2.83
2.56
2.49
2.35
2.41
2.38
2.36
2.35
2.35
2.37
2.35
2.36
2.36
2.36
2.37
2.35
0.29
0.44
0.59
0.74
0.89
1.04
1.18
1.33
1.48
1.63
1.78
1.93
2.07
2.22
2.37
2.52
2.67
5.45
5.49
5.54
5.54
5.58
5.56
5.58
5.55
5.48
5.44
5.38
5.33
5.34
5.26
5.30
5.18
5.23
1.99
1.99
1.98
1.98
1.99
1.99
2.00
2.00
2.00
2.03
2.05
2.01
2.02
2.02
2.04
2.02
2.02
Note: x = 2, logβ = 5.34 ± 0.15, and logk = –2.01 ± 0.03 (25.0°C).
C
dependence was plotted (Fig. 2, line 5). Since the effec-
plex with R = Citr and x = 2 (Fig. 2, line 3) is much
lower than the corresponding value logβ = 18.39 ±
tive rate constant for the redox decomposition of the
complex k , l/(mol s), depends on the extinction coef-
obs
0
.18 for the complex of cerium(IV) with malic acid
ficient ε (k = ε lk , where l is the thickness of the
C
obs
C
C
(
R = Malt, x = 3, Fig. 2, line 4) [5]. The multiple-charge
absorbing solution layer), we also calculated the loga-
3
+
CeOH ion displaces the alcohol hydroxyl proton
along with protons of two carboxyl groups from the
malic acid molecule and forms a complex with the
malate ion much more stable than the complex with the
citrate ion. The difference in the structure and stability
of cerium(IV)–citrate and cerium(IV)–malate interme-
diate complexes can be the reason for the well known
difference between the oscillation regimes of the BZh
reaction for citric and malic acids [30].
–1
rithm of the true rate constant k , s , for the redox
C
decomposition of the complex,
logk = logk – logε l.
(10)
C
obs
C
The extinction coefficient of the complex was deter-
mined by the Beneshi–Hildebrand method [42].
It follows from an analysis of the dependence of
The rate constant for the redox decomposition of
0
0
C
˙
+
–D on the concentration of the complex c
(Fig. 2,
[
(CeOH)H Citr] complexes was calculated from the
−2
0
line 5) and the constant value of logk calculated from
˙
D , D – c /c data series (table) using the equilib-
R M
0
C
–
0
˙
0
rium concentration of the complex at time τ = 0 calcu- the –D , D – c /c data series (table) that one interme-
R
M
0
lated from the D – c /c data series. The logarithm of diate complex predominates in the system. The equality
R
M
the observed rate constant for the intramolecular redox to zero of the free term of Eq. (9) can be evidence of the
decomposition of the complex was analytically calcu- absence of a bimolecular reaction contribution to the
lated at each point of the dependence
redox process (see Fig. 2, line 5). It follows that the law
of the rate of the observed redox reaction is
0
0
C
˙
logkobs = logD – logc
(8)
0
0
˙
–
D = ε lk c ,
(11)
C
C C
with subsequent averaging of the logkobs values
(
table). The k value was also estimated graphically by
obs
and the intermediate mechanism is the mechanism of
the process.
0
˙
0
the –D , D – c /c method. For this purpose, the
R
M
0
0
C
Taking (4) and (7) into account, Eq. (11) can be
rewritten as
˙
–
D = A + k c
(9)
obs
RUSSIAN JOURNAL OF PHYSICAL CHEMISTRY A Vol. 83 No. 6 2009

THE KINETICS OF CERIUM(IV) SULFATE REACTION
949
logkC
dc
dt
–2
M
3+
+
–
------- -- = k K[CeOH ][Citr][H ]
C
(12)
1
3
+
2–
–
–
2
4
4
5
II
=
k β[CeOH ][H Citr ].
C –2
2
Equation (12) establishes a functional dependence of
the rate on the equilibrium concentration of hydrogen
ions in solution. It is in agreement with the general form
of the rate equation –d[Ce(IV)]/dt = k'[Ce(IV)][Citr]
3
I
[
22] and the inverse quadratic dependence of the rate
1
0
20
30 logβ
pseudoconstant k' on the equilibrium concentration of
hydrogen ions in solution also obtained in [22]. Accord-
ing to (12), the latter dependence can be described by the
Fig. 3. logk – logβ correlation dependences: I, for com-
C
+
2
+
equation k' = k K/[H ] . The form of Eq. (12) shows that
plexes of the type [(CeOH)H R] , R = (1) Citr and Ox,
2) R = Lact and Quin, and (3) R = Alk , and II, for com-
plexes of the type [(CeOH)H–xR]3 – x, x > 2, (4) R = Malt
and (5) R = Tart.
C
–2
(
the form of the reagents participating in the elementary
n
2
–
n + 1
3+
redox event is R = H Citr , M
= CeOH rather than
–2
2
2
+
R = Citr, Mn + 1 = Ce(OH) as suggested in [22]. This
modifies the reaction scheme suggested in [22] for the
initial redox process stages as follows:
hols (R = Alk , n = 3, 6 , x = 2, logβ ~ 22.4) [8] and the
n
–
logβ , and logk values obtained for the cerium(IV)–
C
CH COOH
CH COO
2
2
CeOH3+
–
citrate complex allow us to establish an inverse linear
correlation between the logarithms of stability con-
stants and logarithms of the rate constants for the redox
3
+
K
COHCOOH + CeOH
COHCOO
(
13)
CH COOH
CH COOH
2
2
+
decomposition of the [(CeOH)H R] complexes with
–2
+
kc
3+
+
+
2H
PRODUCTS + Ce + 2H .
dibasic ligands of the type R = H L = Citr, Ox; R =
2
H(OH)L = Lact, Quin; R = L(OH) = Alk , n = 3, 6
2
n
The logk = 1.99 ± 0.03 (25°C) value found for the
C
(
Fig. 3, I). This correlation demonstrates the stabilizing
+
[
(CeOH)H Citr] complex coincides in the order of
−2
action of the oxy groups of ligands in complexes with
3
magnitude with the logarithm of the rate constant for the predominantly ionic character of bonds. An
the oxidation of malonic [32] and oxalic [38] dicarbox- increase in the basicity of oxycarboxylic acids above
ylic acids that do not contain oxy groups in molecules two (x > 2) and, as a consequence, an increase in the
with cerium(IV), which substantiates the participation thermodynamic stability of complexes does not influ-
of just carboxyl groups in the formation of the cerium– ence their reactivity (Fig. 3, II). The correlation equa-
citrate complex. This leads us to suggest a similar struc-
ture for the complex of cerium(IV) with malonic acid.
tion for x = 2 has the form logk = –(1.3 + 0.1logβ ).
C
It should, however, be used with care because of the sta-
This value is higher than logkC = 3.40 ± 0.04 for tistical spread of values.
cerium(IV) complexes with α-oxycarboxylic acids
The results obtained in this work lead us to suggest
formed with the displacement of the alcohol hydroxyl
the existence of the corresponding correlations between
proton [5] and is larger by two orders of magnitude than
the electrochemical and optical characteristics of
3
– x
the logkC = 4.2 ± 0.1 value for the corresponding [(CeOH)H R]
complexes [27]. Studies of such cor-
–x
+
relations are currently at the initial stage but are well
worth the effort.
[
(CeOH)H R] complexes with aliphatic polyhydric
−2
alcohols (R = Alk , n = 3, 6 ) [8]. A comparison of the
n
kinetic data on this series of organic substances addi-
tionally substantiates the suggestion that, as distinct
from the α-oxycarboxylic acids considered earlier
REFERENCES
1
2
3
. P. O. Kvernberg, E. W. Hansen, B. Pedersen, et al.,
J. Phys. Chem. A 101, 2327 (1997).
(
malic, tartaric, lactic, etc.), citric acid interacts with
cerium(IV) as a dicarboxylic acid.
. B. Neumann, O. Steinbock, S. C. Muller, et al., J. Phys.
Chem. A 101, 2743 (1997).
The rate constants for the formation of intermediate
cerium(IV) complexes with oxalic (R = Ox, x = 2,
logβ ~ 7.0) [38], quinic (R = Quin, x = 2, logβ ~ 15.0),
lactic (R = Lact, x = 2, logβ ~ 15.0), malic (R = Malt,
x = 3, logβ ~ 18.4), and tartaric (R = Tart, x = 4,
logβ ~ 27.8) carboxylic acids [5] and polyhydric alco-
. K. Butler, O. Steinbock, B. Steinbock, et al., Int. J.
Chem. Kinetics 30, 899 (1998).
3
For complexes with the predominantly covalent type of metal–
n + 1
3+
ligand bonds (M
observed [25].
= Mn ), the opposite dependence is
RUSSIAN JOURNAL OF PHYSICAL CHEMISTRY A Vol. 83 No. 6 2009

9
50
VOSKRESENSKAYA, SKORIK
4
5
6
7
8
9
. Y.-O. Yu and J.-J. Jwo, J. Chin. Chem. Soc. 47, 433 25. Coordination Chemistry of Rare-Earth Elements, Ed. by
(2000).
V. I. Spitsyn and I. Martynenko (Mosk. Gos. Univ., Mos-
cow, 1979).
. O. Voskresenskaya and N. A. Skorik, Russ. J. Appl.
Chem. 75, 820 (2002).
26. Yu. M. Kiselev and N. A. Dobrynina, Coordination
Chemistry (Academie, Moscow, 2007) [in Russian].
. K. Pelle, M. Wittmann, Z. Noszticzius, et al., J. Phys.
Chem. A 107, 2039 (2003).
. K. Pelle, M. Wittmann, K. Lovric, et al., J. Phys. Chem.
A 108, 5377 (2004).
. O. Voskresenskaya and N. A. Skorik, Russ. J. Inorg.
Chem. 50, 1089 (2005).
. S. G. Sobel, H. M. Hastings, and R. G. Field, J. Phys.
Chem. A 110, 5 (2006).
2
2
2
7. Yu. M. Kiselev, Russ. J. Inorg. Chem. 47, 473 (2002).
8. Yu. M. Kiselev, Russ. J. Inorg. Chem. 52, 1826 (2007).
9. B. P. Belousov, Oscillations and Traveling Waves in
Chemical Systems, Ed. by R. J. Field and M. Burger
(Wiley, New York, 1985), p. 605.
3
0. A. M. Zhabotinsky, Oscillations and Traveling Waves in
Chemical Systems, Ed. by R. J. Field and M. Burger
1
1
1
1
1
1
0. S. Balaji, S. Chung, and T. Ramesh, Chem. Eng. J. 126,
1 (2007).
1. L. Van de Water, S. Bulcock, A. Masters, et al., Ind. Eng.
(
Wiley, New York, 1985), p. 615.
1. G. T. Kasperek and T. C. Bruice, Inorg. Chem. 10, 382
1971).
5
3
3
3
3
3
(
Chem. Res. 46, 4221 (2007).
2. R. J. Field, E. Körös, and R. M. Noyes, J. Am. Chem.
Soc. 94, 8649 (1972).
2. J. W. Lee, S. J. Chung, and S. Balji, Chemosphere 68,
1
067 (2007).
3. R. J. Field and R. M. Noyes, J. Chem. Phys. 60, 1877
3. S. Balaji, V. Kokovkin, and S. Chung, Water. Res. 41,
423 (2007).
4. Y. Yamamoto, M. Mori, Y. Aiba, et al., Nucl. Acids Res.
5, e53 (2007).
(1974).
1
4. L. Györgyi, T. Turányi, and R. J. Field, J. Phys. Chem.
4, 7162 (1990).
9
3
5. A. F. Taylor, Progr. React. Kinet. Mech. 27, 247 (2002).
5. R. Thiruvenkatachari, T. Kwon, J. Jun, et al., J. Hazard.
Mater 142, 308 (2007).
36. V. V. Vashin and A. C. Balazs, J. Chem. Phys. A 126,
24707 (2007).
1
1
1
6. H. Yao, X. Yang, and H. Li, Anal. Lett. 39, 2006 (2007).
7. J.-Y. Luo, M. Meng, and Y. Quan, Catal. Lett. 116, 50
2007).
8. S. M. Alam and S. M. Karim, Appl. Chem. 11, 141
2007).
3
7. L. Sciastia, R. Lombardo, and M. L. Turko Liveri,
J. Phys. Chem. B 111, 1354 (2007).
(
3
8. O. Voskresenskaya and N. A. Skorik, in Proc. of XVIIth
Intern. Chug. Conf. on Coord. Chemistry (Minsk, 1990),
Vol. II, p. 341.
1
1
2
2
2
2
2
(
9. M. Getsova, D. Todorovsky, and V. Enchev, Monatshefte
Chem. 138, 389 (2007).
39. I. M. Kolthof, Volumetric Analysis (Intersci., New York,
957), Vol. 3.
1
0. E. T. Denisov, Rate Constants for Homolytic Liquid-
Phase Reactions (Nauka, Moscow, 1971) [in Russian].
1. K. K. Sengupta and S. Aditya, Z. Phys. Chem. (N.F.) 38,
40. L. G. Sillen and A. E. Martell, Stability Constants of
Metal-Ion Complexes, Special Publ. No. 17 (Chem.
Soc., London, 1964).
2
5 (1963).
4
1. L. V. Trubacheva and N. I. Pechurova, Russ. J. Inorg.
2. R. N. Mehrotra and S. Ghosh, Z. Phys. Chem. (DDR)
24, 57 (1963).
3. D. Nebel and G. Urban, Z. Phys. Chem. (DDR) 233, 73 42. N. A. Kostromina, V. N. Kumok, and N. A. Skorik, Coor-
1966).
dination Chemistry (Vyssh. Shkola, Moscow, 1990).
Chem. 26, 3254 (1981).
2
(
4. N. I. Pechurova, G. P. Vakhramova, and V. I. Spitsyn, 43. K. Bachmann and K. H. Lieser, Ber. Bunsges. Phys.
Russ. J. Inorg. Chem. 19, 2074 (1974).
Chem. 67, 810 (1963).
RUSSIAN JOURNAL OF PHYSICAL CHEMISTRY A Vol. 83 No. 6 2009