Unexpected phenomena in the mercury(II)-chlorite ion system: Formation and kinetic role of the HgClO2+ complex

Article abstract of DOI:10.1021/jp0015832

On the basis of time-resolved spectrophotometric measurements, the formation of a novel chlorito complex, HgClO₂⁺, was confirmed under acidic condition in the mercury(II)-chlorite ion system. The complex formation occurs within the dead time of the stopped-flow instrument applied (τ_d approximately 1 ms). The stability constant of HgClO₂⁺ was determined from the initial part of the kinetic traces recorded at the characteristic 325 nm absorbance maximum of the complex: K = 41.7±1.9 M^-1, ε₃₂₅ = 727±24 M^-1 cm^-1 (I = 1.0 M (NaClO₄), 25.0 °C). An intense formation of chlorine dioxide was observed which is due to mercury(II)-catalyzed decomposition of chlorite ion. At longer reaction times, when most of the chlorite ion is consumed, the concentration of chlorine dioxide abruptly drops to its final value. This unique feature of the reaction is attributed to an autocatalytic redox cycle which is initiated by an electron transfer step between the coordinated chlorite ion and chlorine dioxide. The proposed mechanism postulates the formation of a Cl(II) intermediate which rapidly oxidizes both ClO₂^- and ClO₂.

Full text of DOI:10.1021/jp0015832

J. Phys. Chem. A 2000, 104, 8045-8049
8045
Unexpected Phenomena in the Mercury(II)-Chlorite Ion System: Formation and Kinetic
+
Role of the HgClO2 Complex
Istva´n Fa´bia´n* and Do´ra Sz_ucs
Department of Inorganic and Analytical Chemistry, UniVersity of Debrecen, Debrecen, Hungary
Gilbert Gordon
Department of Chemistry and Biochemistry, Miami UniVersity, Oxford, Ohio 45056
ReceiVed: April 25, 2000; In Final Form: June 16, 2000
On the basis of time-resolved spectrophotometric measurements, the formation of a novel chlorito complex,
+
HgClO2 , was confirmed under acidic condition in the mercury(II)-chlorite ion system. The complex formation
occurs within the dead time of the stopped-flow instrument applied (τd ∼ 1 ms). The stability constant of
+
HgClO2 was determined from the initial part of the kinetic traces recorded at the characteristic 325 nm
absorbance maximum of the complex: K ) 41.7 ( 1.9 M-1, ꢀ325 ) 727 ( 24 M-1 cm-1 (I ) 1.0 M (NaClO4),
25.0 °C). An intense formation of chlorine dioxide was observed which is due to mercury(II)-catalyzed
decomposition of chlorite ion. At longer reaction times, when most of the chlorite ion is consumed, the
concentration of chlorine dioxide abruptly drops to its final value. This unique feature of the reaction is
attributed to an autocatalytic redox cycle which is initiated by an electron transfer step between the coordinated
chlorite ion and chlorine dioxide. The proposed mechanism postulates the formation of a Cl(II) intermediate
-
which rapidly oxidizes both ClO2 and ClO2.
Introduction
unexpected phenomenon and indicates very unique features of
the Hg(II)-ClO2--ClO2 system.
Chlorite ion1 is well-known for its complex redox chemistry
in acidic aqueous solution. It was shown before that different
types of interactions may exist between chlorine(III) and metal
ions.2-15 Besides a kinetically inert complex,7,15 [(NH3)5Co-
(ClO2)]2+, the formation of labile complexes was confirmed with
U(VI),2 Cu(II),11 and Fe(III).12 The uranyl complex is too weak
and its properties could not be quantified. The stability constants
The observations imply that ClO2 is involved in some sort
-
of redox process and either oxidized to ClO3 or reduced to a
lower oxidation state chlorine species. However, a direct
involvement of Hg(II) in such a process can be excluded because
it is not a strong enough oxidant to oxidize ClO2 and it cannot
be oxidized to a +III oxidation state either. This leads to the
conclusion that the overall process proceeds through the
formation of a so far unidentified intermediate which opens an
+
and molar absorptivities of the other two species,16 CuClO2
and FeClO22+, were determined on the basis of stopped-flow
studies. While Cu(ClO2)+ is inert toward redox decomposition,11
the formation of FeClO22+ is the key reaction step in the redox
cycle which leads to catalytic decomposition of chlorine(III).12,13
The oxidation of low oxidation state metal ions by chlorine-
(III) typically shows complex stoichiometric and kinetic
patterns.3-6,8 Though chlorito complexes were not detected in
these systems, it was assumed that some sort of association exists
between the reactants and the reactions proceed via inner-sphere
redox mechanisms.
efficient reaction path to remove ClO2. In the present study,
+
we report the formation of a novel chlorito complex, HgClO2
,
and show that this species plays a significant role in the
consumption of chlorine dioxide.
Experimental Section
Reagents. Acidic mercury(II) stock solution was prepared
by adding analytical grade HgO (REANAL, Hungary) to 70%
perchloric acid (Carlo Erba) in excess. The sludge was filtered
after extensive stirring for 3-4 h. Excess liquid was removed
by heating the mother liquor on a water bath and mercury(II)
perchlorate was allowed to precipitate. The crystals were filtered
off, washed with ethanol, and dissolved in dilute perchloric acid.
The stock solution was purged with ozone for a brief period of
time to remove any trace of Hg(I). Tests by adding concentrated
NaCl solution to an aliquot of the stock solution were negative
for Hg(I). Kinetic runs were identical regardless whether the
samples were pretreated with ozone or not. Mercury(II)
concentration was determined with an ICP-AES method by
using a Spectroflame ICP instrument (Spectro GmbH, Ger-
many). The acid content was determined by pH-metric titration
with standardized NaOH solution, first converting Hg(II) into
tetraiodo complex by adding excess KI to the sample. Sodium
Redox reactions of chlorite ion are often coupled with
catalytic decomposition of this species producing chlorine
dioxide which is stable under acidic condition.17 As an excep-
tion, the disappearance of ClO2 was found in acidic aqueous
solution of chlorite ion, chlorine dioxide, and mercury(II).18 This
phenomenon was used to confirm the formation of chlorous
acid as an intermediate in the oxidation of methanol with
chlorate ion. On the basis of the chemical properties of
oxychlorine species and Hg(II), the decay of ClO2 is quite an
* Corresponding author. Address: Department of Inorganic and Analyti-
cal Chemistry, University of Debrecen, Debrecen P.O.B. 21, Hungary
H-4010. Tel: +36 52 512-900, ext 2378. Fax: +36 52 489-667. E-mail:
ifabian@delfin.klte.hu.
10.1021/jp0015832 CCC: $19.00 © 2000 American Chemical Society
Published on Web 08/04/2000

8046 J. Phys. Chem. A, Vol. 104, No. 34, 2000
Fa´bia´n et al.
Figure 1. Characteristic spectral change as a function of time in the
mercury(II)-chlorite ion reaction. CHg(II)0 ) 0.015 M; CCl(III)0 ) 0.010
M; [H+] ) 0.025 M; 25.0 °C; I ) 1.0 M (NaClO4).
Figure 2. Stopped-flow traces as a function of acidity in the mercury-
0
0
(II)-chlorite ion reaction. CHg(II) ) 0.030 M; CCl(III) ) 0.0050 M;
[H+] ) 0.350, 0.300, 0.250, 0.200, 0.180, 0.160, 0.140, 0.120, 0.100
M in the order of increasing absorbance at the maximum. 25.0 °C, I )
1.0 M (NaClO4).
chlorite was obtained by purifying commercially available
NaClO2 (Fluka, 80% purity) as described earlier.11 All other
reagents were of the highest available purity. Reagent solutions
and samples were prepared in doubly deionized and ultrafiltered
water obtained from a MILLI-Q RG (Millipore) water purifica-
tion system. Measurements were made at 25.0 ( 0.1 °C. The
ionic strength was adjusted to 1.0 M with NaClO4 prepared from
perchloric acid and Na2CO3 (Reanal, Hungary).
Instrumentation and Methods. Potentiometric titrations
were made with a Metrohm 721 NET Titrino system. Time-
resolved spectral changes were monitored with either a HP-
8543 diode array spectrophotometer equipped with an APL
(Applied Photophysics) RX 2000 rapid mixer or an APL DX-
17 MV sequential stopped-flow apparatus at 10 mm optical path.
In the latter case, the spectra were generated from individual
stopped-flow traces obtained at each 10 nm in the wavelength
region studied by using the software package provided with the
instrument. Fitting of the experimental data and kinetic model
calculations was done with the program packages PSEQUAD19
and SCIENTIST,20 respectively.
reaction mixture. Apparently, the actual course of the reaction
strongly depends on the acidity, the concentrations, and
concentration ratios of the reagents. When [H+] was increased,
the transient amount of ClO2 was decreased and the reaction
became faster. However, as it was confirmed by stopped flow
measurements, the kinetic pattern was basically the same under
all conditions applied (Figure 2). The formation of chlorine
dioxide is partly due to acid-catalyzed decomposition of
chlorine(III) in this system. The stoichiometry and kinetics of
the corresponding reactions are well established and were
reported earlier.21 Mercury(II) also catalyzes the decomposition
because the initial absorbance change at the characteristic ClO2
band became much faster by addition of Hg(II) to the samples.
Subsequent decay of the absorbance implies that chlorine
dioxide, which is otherwise inactive in acidic solution and in
the absence of reducing agents, is an intermediate in this system.
An inspection of Figure 1 reveals that the absorbance
maximum was shifted to higher wavelengths during the reaction
until it reached 360 nm. Immediately after mixing the reactants,
the absorbance maximum was observed at 325 nm and the
characteristic fine structure of ClO2 was not seen in the spectra.
The 325 nm band cannot be assigned to Hg(II), ClO2-, or ClO2,
and the observations clearly confirm the formation of a novel
species in the initial phase of the reaction. According to stopped-
flow measurements at 400 nm, chlorine dioxide formation is
relatively slow and does not interfere with spectral measure-
ments in the first few hundred milliseconds. At longer reaction
times the absorbance of ClO2 suppresses other spectral effects.
The formation of the new species, observed as an absorbance
jump at the beginning of the kinetic traces, was too fast to be
monitored by the stopped-flow method (τd ∼ 1 ms). The
concentration and pH dependencies of the initial absorbance
change were studied at 325 nm by using the average absorbance
value from the first 5-100 ms section of each stopped-flow
trace. Evaluation of the data by using the Job method22
confirmed the formation of a new species with a 1:1 stoichi-
ometry for ClO2- and Hg(II). We assigned the spectral effects
to the formation of the mercury(II)-chlorito complex. Ligand
substitution reactions of mercury(II) are typically diffusion
controlled, and the extremely fast initial absorbance change lends
support to this assumption. The data were evaluated on the basis
of the following equilibrium model:
Results and Discussion
In the first set of experiments, the spectra of various reaction
mixtures of mercury(II), chlorite ion and chlorine dioxide were
measured as a function of time by using conventional spectro-
photometry. In order to prevent the precipitation of mercury
hydroxide, all samples were acidified. The spectra of chlorine
dioxide solutions containing mercury(II) did not change in a
sealed spectrophotometric cell for several hours. Chlorine
dioxide was also stable when chloride ion was added to the
samples in 0.001-0.1 M concentration. When chlorite ion was
added to a solution of mercury(II) and chlorine dioxide, a
significant absorbance decay was observed at the characteristic
absorbance band of ClO2 (λmax ) 360 nm)17 after an incubation
period.
Upon mixing acidic solutions of mercury(II) and chlorite ion,
first an intense absorbance increase was observed. The absor-
bance abruptly dropped to its final value at longer reaction times
(Figure 1). One of the dominant absorbing species is chlorine
dioxide in the near-UV-visible spectral region and, as a first
approach, the observations correspond to the formation and
subsequent decay of this species. The fine structure of the spectra
and the absorbance maximum of the final spectrum at around
360 nm corroborates this assumption. Under certain conditions
all chlorine dioxide, produced in the initial phase, was depleted
and the reaction mixture became colorless. In other cases, the
decay did not go to completion and some ClO2 remained in the
[HClO2]
H+ + ClO2- h HClO2
Kp )
(1)
[H+][ClO2 ]
-

Mercury(II)-Chlorite Ion System
J. Phys. Chem. A, Vol. 104, No. 34, 2000 8047
+
[HgClO2 ]
Hg2+ + ClO2- h HgClO2
K )
(2)
+
[Hg2+][ClO2 ]
-
The contribution of mercury(II), chlorite ion and chlorous
acid to the measured absorbance is negligible at 325 nm.
Provided that the acidity is constant, mercury(II) is in large
excess over chlorite ion and the chlorito complex is the only
absorbing species in the initial phase, the following expression
applies:
K′CHg(II)
+1 + K′CHg(II)
+
A ) ꢀHgClO +[HgClO2 ] ) ꢀ
CCl(III) (3)
HgClO2
2
-
where ꢀHgClO , K′, CHg(II), and CCl(III) are the molar absorptivity
2
Figure 3. Plot of the absorbance according to eq 4 as a function of
acidity at constant chlorite ion concentration. CCl(III)0 ) 0.0020 M; [H+]
) 0.006 M (b), 0.010 M (O), 0.020 M (2), 0.035 M (4), 0.040 M
(9), 0.060 M (0); 25.0 °C, I ) 1.0 M (NaClO4).
of the complex, the pH-dependent apparent stability constant,
and the total initial concentrations of mercury(II) and chlorite
ion, respectively. After rearrangement, eq 3 yields
TABLE 1: Equilibrium Constants for the Reactions of
1
A
1
1
1
-
)
+
(4)
ClO2 (1.0 M NaClO4; 25.0 °C)
ꢀHgClO +CCl(III) K′ꢀHgClO +CCl(III) CHg(II)
2
2
reaction
K
ref
H+ + ClO2- ) HClO2
52.5
36.3
1.04
13.8
13.2
41.7 ( 1.9
11
The data are plotted according to eq 4 in Figure 3. In the
final calculations, Kp was fixed at the value reported earlier,11
log Kp ) 1.72, and all experimental data were fitted simulta-
14a
11
Cu2+ + ClO2- ) CuClO2
+
Fe3+ + ClO2- ) FeClO2
12
2+
neously on the basis of eq 3. The results are K ) 41.7 ( 1.9
14a
M-1 and ꢀHgClO (at 325 nm) ) 727 ( 24 M-1 cm-1
.
+
Hg2+ + ClO2- ) HgClO2
a At 5.0 °C.
this work
+
2
Available stability constants for the chlorito complexes are
collected in Table 1. As shown, mercury(II) forms about 3 and
40 times more stable complex than iron(III) and copper(II),
respectively. The interpretation of these differences is not
straightforward. An essential question is whether chlorite ion
is coordinated via oxygen or chlorine atoms to the metal center.
Structural data or, in the case of iron(III), a well-defined
spectrum is not available for chlorito complexes because redox
decomposition and coupled side reactions prevent thorough
characterization of these species. Thus, the considerations
regarding the coordination mode(s) of ClO2- need to be based
on the comparison of the equilibrium constants.
+
(II)]/[Cl(III)] ratio. Concentration profiles for the HgClO2
complex and ClO2 were calculated from time-resolved spectra
as shown in Figure 5. A comparison of these kinetic traces
reveals that most likely the chlorito complex controls the
reaction sequence which leads to the decay of chlorine dioxide.
At constant CHg(II) and pH, the actual concentration of HgClO2
is determined by CCl(III). Thus, after instantaneous formation of
+
+
the complex, the concentrations of CCl(III) and [HgClO2
]
decrease simultaneously. When chlorite ion is used up, the
chlorito complex disappears and, as consequence, the depletion
of chlorine dioxide also ceases. We propose the following kinetic
model for the interpretation of the results:
Provided that O-coordination occurs, simple electrostatic
considerations would predict the Fe(III) > Cu(II) ∼ Hg(II)
stability order. The deviation from this trend is a strong
indication that other factors are also important. The “irregular”
stability order can be justified by taking into account the hard-
soft character of the metal ions. A specific interaction between
the soft mercury(II) ion and the relatively soft chlorine atom
may stabilize the HgClO2+ complex. Such a preference of Hg-
(II) toward a soft donor atom is well established in the case of
mercury(II)-sulfito complexes.23 In contrast, the soft-soft
stabilization effect is very unlikely with iron(III) and copper-
(II) which are typical hard metal ions. Thus, the coordination
mode of the ligand is not necessarily uniform in all chlorito
complexes. The Cl-coordination seems to be favorable in the
case of the soft mercury(II) while the O-coordination is expected
to prevail in the other two complexes. These considerations are
fully supported by the X-ray structure of the inert [(NH3)5Co-
(ClO2)][Pd(CN)4]‚H2O complex in which chlorite ion is coor-
dinated to the hard cobalt(III) center via an oxygen atom.15
The kinetic profiles show distinct variation as a function of
H+ + ClO2- h HClO2
Hg2+ + ClO2- h HgClO2
fast
fast
(R1)
(R2)
+
ClO2- f ClO2 + ClO3- + Cl-
slow initiation (R3)
HgClO2+ + ClO2 f Hg2+ + ClO3- + Cl(II)
rate det
(R4)
ClO2- + Cl(II) + H+ f ClO2 + HOCl
ClO2 + Cl(II) f ClO3- + HOCl
fast (R5)
fast
(R6)
ClO2- + HOCl f ClO2 + Cl- (+ ClO3 )
fast (R7)
-
-
reagent concentration and pH. At relatively high ClO2 con-
The concentration of the HgClO2+ complex, a key intermedi-
ate in this system, is controlled by (R1) and (R2) which are
fast preequilibria. According to the model, chlorine dioxide is
initially produced by relatively slow acid-catalyzed decomposi-
tion of chlorite ion, (R3). This reaction triggers the autocatalytic
cycle and has marginal kinetic importance at longer reaction
centration, the stopped-flow traces at 360 nm feature a plateau
before the absorbance decreases to its final value (Figure 4). In
excess Hg(II), the absorbance increases at a growing rate in
the initial phase indicating autocatalytic generation of ClO2
(Figure 4, inset). This feature disappears on decreasing the [Hg-

8048 J. Phys. Chem. A, Vol. 104, No. 34, 2000
Fa´bia´n et al.
and the main products are chloride ion and chlorine dioxide
when chlorite ion is present in excess over hypochlorous
acid.26,27
In model calculations, kinetic traces were simulated by
assuming the simplest rate expressions for the individual steps
and ignoring the pH dependence of the reaction. When an
appropriate set of rate constants was selected, the calculations
predicted that the ClO2 concentration goes through a maximum
as a function of time. The noted autoinhibition could also be
reproduced and, similarly to the experimental observations, the
maximum amount of ClO2 increased with increasing [ClO2-].
The model appears to be less satisfactory for reproducing the
exact shape of the kinetic curves. (The kinetic model and
simulated kinetic traces are given in the Supporting Information.)
Quantitative evaluation of the mechanism would require
thorough kinetic description of the component reactions. The
complexity of reaction R7 is well documented and its kinetics
and stoichiometry have been extensively studied before.26,27
However, earlier proposed mechanisms are valid only for the
slightly acidic neutral pH region and cannot be used for the
conditions applied in the present study. It is open to question
whether sufficient experimental information can be collected
for the concentration and pH dependencies of the reaction rates
and stoichiometries of steps R5, R6, and R7. Further complica-
tions may arise due to the formation of chloride ion. It forms
relatively stable complexes with mercury(II) and may deactivate
the catalyst.
Figure 4. Stopped-flow traces as a function of chlorite ion concentra-
tion in the mercury(II)-chlorite ion reaction. CHg(II)0 ) 0.030 M; [H+]
) 0.350 M; CCl(III)0 ) 0.0020 M (1), 0.0050 M (2), 0.010 M (3), 0.015
M (4), 0.020 M (5); 25.0 °C, I ) 1.0 M (NaClO4). Inset: Enlarged
initial sections of traces 1 and 2.
In spite of the limitations, the proposed model provides
appropriate qualitative interpretation of the results and consistent
with the experimental observations. The autocatalytic nature of
the overall reaction can be interpreted on the basis of the (R4)-
(R6) reaction sequence. In these reaction steps, more ClO2 is
formed than consumed and (R4) becomes faster as the reaction
proceeds. Because [HgClO2+] is proportional to the total
concentration of chlorite ion, eventually the autocatalytic effect
+
Figure 5. Kinetic profiles for HgClO2 (b) and ClO2 (2) in the
mercury(II)-chlorite ion reaction. CHg(II)0 ) 0.015 M; CCl(III)0 ) 0.010
M; [H+] ) 0.025 M; 25.0 °C, I ) 1.0 M (NaClO4).
-
is offset by the consumption of ClO2 in step R7. The abrupt
decay of ClO2 at the end of the reaction implies that Cl(II)
oxidizes chlorine dioxide much faster than chlorite ion. Thus,
as [ClO2-] decreases the consumption of ClO2 in step R4
becomes superior compared to its production in (R5) and (R7).
Provided that sufficient kinetic information becomes available
for the component reactions, the results presented here can serve
as a basis to develop a detailed mechanism for the Hg(II)-
ClO2--ClO2 reaction.
times. The consumption of ClO2 occurs only when mercury(II)
and chlorite ion are simultaneously present in the reaction
mixture. This strongly suggests that chlorine dioxide reacts
directly with the chlorito complex in the rate-determining step,
(R4). Precipitation was not observed either during the reaction
or when excess chloride ion was added to the reaction mixtures
confirming that mercury(I) does not form in the redox process.
It follows that direct involvement of Hg(II) in any redox step
can be excluded and the electron transfer probably occurs
between the coordinated chlorite ion and ClO2. The model
predicts the formation of a Cl(II) intermediate in step R4. The
formation of Cl(II) was proposed in a few redox reactions of
chlorite ion before,4,8,13 but direct experimental evidence is not
available to confirm the exact composition of this species. It
Acknowledgment. This work was supported by the Hungar-
ian National Research Foundation under grant No. OTKA T
029568 and M 028244 and by a NATO Linkage Grant under
grant No. CRG.LG 973337.
Supporting Information Available: The simplified kinetic
model and a figure with simulated kinetic traces. This material
is available free of charge via the Internet at http://pubs.acs.org.
-
can be ClO which is expected to react very rapidly with ClO2
(R5) and/or with other species in the system. According to pulse
radiolysis studies, the lifetime of ClO is on the order of several
References and Notes
10 µs in aqueous solution and the second-order decay of ClO
- 24 The product of this reaction is ClO3
,
-
(1) Chlorite ion and chlorous acid are in fast equilibrium, and their
concentration ratio is determined by the pH. The two species will not be
distinguished in this paper and will be referred to as either chlorite ion or
chlorine(III).
(2) Gordon, G.; Kern, D. M. Inorg. Chem. 1964, 3, 1055.
(3) Thompson, R. C.; Gordon, G. Inorg. Chem. 1966, 5, 562.
(4) Ondrus, M. G.; Gordon, G. Inorg. Chem. 1972, 11, 985.
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1981, 2537.
is catalyzed by ClO2
.
implying that ClO is also capable of oxidizing chlorine dioxide
in a fast reaction step, (R6). The shapes of the kinetic traces
are determined by the competition of chlorite ion and chlorine
dioxide for the same reactive intermediate The sharp decay of
ClO2 at the end of the reaction strongly suggests that (R6) is
faster than (R5). Such a competition between ClO2- and ClO2
was postulated in the mechanism for the reduction of bromate
ion by chlorous acid.25 Under acidic condition, step R7 is fast

Mercury(II)-Chlorite Ion System
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