Free-radical-induced oxidation of hydroxymethanesulfonate in aqueous solution: Part 2. A pulse and steady-state radiolysis study of oxygenated solutions

Article abstract of DOI:10.1039/a703266h

At pH 4.0, the radical ^.CH(OH)SO₃^- formed by reaction of ^.OH with hydroxymethanesulfonate (HMS) reacts rapidly with oxygen to form the peroxyl radical, ^.OOCH(OH)SO₃^-, with k₄ = (2.6 ± 0.3) × 10⁹ dm³ mol^-1 s^-1. This radical decays by a first-order process with k = 2.4 × 10⁴ s^-1 and E_a = (29 ± 1.2) kJ mol^-1. The decay is by a split pathway, reactions (6) and (7), with k₆ = 1.7 × 10⁴ s^-1 and k₇ = 7.0 × 10³ s^-1. Reaction (6) involves direct elimination of HO₂^.. Reaction (7) initially forms the formylperoxyl radical, but this is hydrolysed, reaction (8), to formic acid and HO₂^. with k₈ = 2.46 × 10³ s^-1. ^.CH(OH)SO₃^- + O₂ → ^.OOCH(OH)SO₃^- (4) ^.OOCH(OH)SO₃^- → HO₂^. + CHOSO₃^- (6) ^.OOCH(OH)SO₃^- → ^.OOCHO + HSO₃^- (7) ^.OOCHO + H₂O → ^.OOCH(OH)₂ → HCOOH + HO₂^. (8) The yield of formate is compatible with this mechanism if it is assumed that the product of reaction (6), i.e. CHOSO₃^-, can also yield formic acid. Yields of SO₄^2- have been determined and their origin is considered. At pH 7.0, the radical ^.CH(O^-)SO₃^- reacts with oxygen to yield O₂^.- with k₁₁ = (1.6 ± 0.1) × 10⁹ dm³ mol^-1 s^-1. ^.CH(O^-)SO₃^- + O₂ → CHOSO₃^- + O₂^.- (11).

Full text of DOI:10.1039/a703266h

View Article Online / Journal Homepage / Table of Contents for this issue
Free-radical-induced oxidation of hydroxymethanesulfonate in
aqueous solution
Part 2¤ A pulse and steady-state radiolysis study of oxygenated solutions
Susan Barlow, George V. Buxton, Samantha A. Murray and G. Arthur Salmon*
Cookridge Radiation Research Centre, University of L eeds, Cookridge Hospital, L eeds,
UK L S16 6QB
At pH 4.0, the radical ~CH(OH)SO ~ formed by reaction of ~OH with hydroxymethanesulfonate (HMS) reacts rapidly with
3
oxygen to form the peroxyl radical, ~OOCH(OH)SO ~, with k \ (2.6 ^ 0.3) ] 109 dm3 mol~1 s~1. This radical decays by a
3
4
Ðrst-order process with k \ 2.4 ] 104 s~1 and E \ (29 ^ 1.2) kJ mol~1. The decay is by a split pathway, reactions (6) and (7),
a
with k \ 1.7 ] 104 s~1 and k \ 7.0 ] 103 s~1. Reaction (6) involves direct elimination of HO ~. Reaction (7) initially forms the
6
7
2
formylperoxyl radical, but this is hydrolysed, reaction (8), to formic acid and HO ~ with k \ 2.46 ] 103 s~1.
2
8
~CH(OH)SO ~ ] O ] ~OOCH(OH)SO ~
(4)
(6)
(7)
(8)
3
2
3
~OOCH(OH)SO ~ ] HO ~ ] CHOSO ~
3
2
3
~OOCH(OH)SO ~ ] ~OOCHO ] HSO ~
3
3
~OOCHO ] H O ] ~OOCH(OH) ] HCOOH ] HO ~
2
2
2
The yield of formate is compatible with this mechanism if it is assumed that the product of reaction (6), i.e. CHOSO ~, can
3
also yield formic acid. Yields of SO 2~ have been determined and their origin is considered.
4
At pH 7.0, the radical ~CH(O~)SO ~ reacts with oxygen to yield O ~~ with k \ (1.6 ^ 0.1) ] 109 dm3 mol~1 s~1.
3
2
11
~CH(O~)SO ~ ] O ] CHOSO ~ ] O ~~
(11)
3
2
3
2
mechanism of the overall oxidation. As outlined in Part 1,13
pulse radiolysis is a convenient means of studying the kinetics
of the processes while the overall products of the oxidation
are readily generated by steady-state radiolysis with a 60Co
c-ray source.
Introduction
Hydroxymethanesulfonate (HMS), which is formed from the
addition of sulÐte/hydrogen sulÐte to formaldehyde, reaction
(1), has been detected in atmospheric droplets1,2 and it has
been proposed that its formation enhances the wet deposition
of formaldehyde and SO from the atmosphere.3 It is also
2
thought that its formation plays an important role in prevent-
Experimental
ing the oxidation of SIV by H O in the aqueous phase of
2
2
clouds.4,5 However, the oxidation of SIV in clouds might also
occur by a free-radical chain process initiated by ~OH radicals
and involving the radicals SO ~~, SO ~~ and SO ~~.6h10
Chemicals
The chemicals used were as described in Part 1.13 In addition,
oxygen (BOC) was taken directly from the cylinder.
3
4
5
Deister et al.11 have concluded that, in the presence of oxygen,
the reaction of ~OH with HMS results in the chain oxidation
of HMS.
Product analysis
HCHO ] HSO ~ H CH (OH)SO ~
(1)
3
2
3
Analysis of formate and sulfate was by ion chromatography.
The system used consisted of a Dionex advanced gradient
pump, a Dionex ionpac AS12A analytical column with an
AG12A guard column, an AAMS-1 suppressor column, a con-
ductivity detector and a chart recorder. The eluent used for
where K \ 3.6 ] 106 mol~1 dm3.12
1
In Part 113 of this series we showed that ~OH reacts with
HMS, reaction (2), to form ~CH(OH)SO ~ and that this
3
radical undergoes acidÈbase dissociation, reaction (3), with a
pK of 6.35 ^ 0.1.
the separation contained 0.3 mmol dm~3 NaHCO and 2.7
a
3
mmol dm~3 Na CO . 10 ml of the solution was irradiated
~OH ] CH (OH)SO ~ ] ~CH(OH)SO ~ ] H O
(2)
(3)
2
3
2
3
3
2
and a 20 ll aliquot injected into the chromatograph by the
complete-Ðlling method. The HCOO~ peak elutes after 2.8
~CH(OH)SO ~ H ~CH(O~)SO ~ ] H`
3
3
min and that of SO 2~ after 19.0 min.
where k \ (3.0 ^ 0.3) ] 108 dm3 mol~1 s~1.
In this paper we present a study of the reactions of
4
2
~CH(OH)SO ~ and ~CH(O~)SO ~ with oxygen and of the
Steady-state irradiation
3
3
Steady-state irradiations were carried out using a 300 Ci 60Co
c-ray source. Solutions were contained in a custom-made cell
¤ Part 1: ref. 13.
J. Chem. Soc., Faraday T rans., 1997, 93(20), 3641È3645
3641

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that allowed the solution to be bubbled in situ with an
N OÈO mixture prior to irradiation. The gas mixture was
2
2
produced by mixing separate Ñows of N O and O , which
2
2
were controlled by needle valves and monitored using Ñow
meters. The resulting solutions contained 6.75 ] 10~4 mol
dm~3 O and 1.25 ] 10~2 mol dm~3 N O. The cell was posi-
tioned reproducibly relative to the c-ray source and the
volume of solution used was always 10 cm3. The mean dose
rate throughout the sample was determined using the Fricke
Dosimeter.14
2
2
Pulse radiolysis
The pulse radiolysis apparatus and procedures when using
spectrophotometric detection have been described in detail
elsewhere.15 Experiments with conductivity detection used the
ac system described by Janata et al.16 The solution was con-
tained in a rectangular Pyrex cell with internal dimensions of
5.0 mm ] 7.0 mm ] 10.0 mm and which had disc-shaped elec-
trodes of platinum foil (3.0 mm diameter) mounted in parallel
conÐguration (separation 5.0 mm) on tungsten wires, which
were along the long axis of the cell. The cell was mounted so
that the electron beam entered the cell perpendicular to the
7.0 mm ] 10.0 mm face. Side-arms in the vertical plane with
B5 cones and sockets allowed the cell to be Ðlled with solu-
tion. A similar cell mounted above the irradiated cell, but
screened from the radiation, acted as a reference.
Fig. 1 Absorption spectra obtained by pulse radiolysis of 0.1 mol
dm~3 HMS solutions at pH 4.0. A, spectrum of ~CH(OH)SO ~
3
obtained in N O-saturated solution; B, spectrum at 2.5 ls; and C,
2
143 ls after the pulse for solutions with [O ] \ 6.75 ] 10~4 mol
2
dm~3.
Rate constant for the reaction. The kinetics of the reaction
were investigated by monitoring the growth of absorption at
245 nm on pulse radiolysis of solutions with [HMS] \ 0.1
mol dm~3 for which the oxygen concentrations varied from
192 to 675 lmol dm~3. The growth of absorption obeyed
For the conductivity measurements, dosimetry was carried
Ðrst-order kinetics and the rate constants obtained, k , were
out using N O-saturated solutions containing 10~3 mol dm~3
obs
proportional to the oxygen concentration with slope, k \ (2.6
4
2
dimethyl sulfoxide at pH 4.4, for which ; (G j ) was taken to
i i
^ 0.3) ] 109 dm3 mol~1 s~1, which is of the order expected
be 2.02 ] 10~8 S m2 J~1,17 where G j is the product of the
radiolytic yield, G , of the ions contributing to the conductiv-
i i
for the addition of oxygen to a carbon-centred radical.
i
ity change in units of mol J~1 and j , their limiting ionic con-
i
Decay
of
~OOCH(OH)SO —.
The
decay
of
ductances in units of S m2 mol~1.
3
~OOCH(OH)SO ~ was monitored at 300 nm for solutions at
Solutions were saturated with argon, nitrous oxide or
oxygen as required for a particular experiment and then
stored in all-glass syringes in readiness for pulse radiolysis. To
prepare solutions containing both N O and O , syringes con-
3
pH 4.0 and was found to be Ðrst-order and independent of
[O ] with a rate constant of (2.4 ^ 0.1) ] 104 s~1. The tem-
2
perature dependence of this rate constant was investigated for
2
2
temperatures between 5 and 40 ¡C. A plot of ln k vs. 1/T was
taining the separate saturated solutions were joined by a cap-
illary tube and then an appropriate volume of one solution
forced into the other. A glass disc in the receiving syringe
allowed the solutions to be mixed by shaking.
linear and gave E \ (29 ^ 1.2) kJ mol~1 and ln(A/
a
s~1) \ 22.3 ^ 0.5.
The spectrum of the product of the decay of
~OOCH(OH)SO ~ is shown in Fig. 1, C. a-Hydroxyperoxyl
Except where otherwise speciÐed, all experiments were
carried out at ambient temperature, 18È20 ¡C.
3
radicals are known to dissociate by elimination of HO ~,18 so
2
the possibility that this is the case for ~OOCH(OH)SO ~ was
3
tested by comparing the spectrum at 143 ls, after allowance
Results and Discussion
for the yield of HO ~ formed from the small yield of H-atom
2
formed on the radiolysis of water [G(H) \ 6 ] 10~8 mol J~1],
Reaction of ~CH(OH)SO — with oxygen
with that expected for HO ~/O ~ at pH 4.0 (see Fig. 2). The
3
2
2
molar absorption coefficients for the experimental spectrum
were calculated assuming that, at the concentration of HMS
Absorption spectrum of ~OOCH(OH)SO —. The reaction of
3
~CH(OH)SO ~ with oxygen was investigated by pulse
3
employed,
G[~OOCH(OH)SO ~] \ G(~OH) \ 6.0 ] 10~7
3
radiolysis with 500 ns, 15 Gy pulses of solutions at pH 4.0
mol J~1. The spectrum of HO ~/O ~~ was calculated from the
2
2
containing 0.1 mol dm~3 HMS and with [N O] \ 1.25
known spectra of HO ~ and O ~~ 19 and K for equilibrium
2
] 10~2 mol dm~3 and [O ] \ 6.75 ] 10~4 mol dm~3. The
2
2
2
a
(5).20
ratio of N O/O in the solutions is such that 90% of the e ~
2
2
aq
generated by radiolysis reacts with N O to yield ~OH.
HO ~ H H` ] O ~~
(5)
2
2
2
At the concentrations of HMS used here reaction (2) is
e†ectively complete during the pulse, and the end-of-pulse
Because of the high ionic strength of the solutions (0.1 mol
dm~3), [H`] and [O ~~] were evaluated using K for equi-
absorption is that of ~CH(OH)SO ~ 13 (see Fig. 1, A). After
2
c
3
librium (5) rather than K . K was estimated from K as indi-
the pulse, a further growth of absorption occurred over about
a
c
a
cated in ref. 21. Also, account was taken of the inÑuence of the
high ionic strength on the glass electrode of the pH meter.
This was done by comparing pH readings for solutions with
[H`] \ 10~4 mol dm~3 in the absence and presence of 0.1
mol dm~3 NaClO . The e†ect of 0.1 mol dm~3 NaClO was
2.5 ls and the absorption then decayed over B100 ls and
then decayed back to the baseline over 2È20 ms. This variabil-
ity was attributed to the inÑuence of trace metal ions on the
decay of HO ~/O ~ (see below). Spectra obtained at 2.5 ls
2
2
and 140 ls after the pulse are also shown in Fig. 1(B and C).
The spectrum observed at 2.5 ls (B) is assigned to the peroxy
radical formed by addition of O to ~CH(OH)SO ~, reaction
4
4
to shift the pH reading by ]0.18.
It is apparent from Fig. 2 that the experimental spectrum
2
3
di†ers from the predicted one for HO ~/O ~ and that
(4).
2
2
~OOCH(OH)SO ~ does not dissociate simply to yield HO ~
3
as in reaction (6).
2
O ] ~CH(OH)SO ~ ] ~OOCH(OH)SO ~
(4)
2
3
3
3642
J. Chem. Soc., Faraday T rans., 1997, V ol. 93

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Fig. 2 Comparison of the spectrum obtained at 143 ls after the
pulse in solutions at pH 4.0 and [O ] \ 6.75 ] 10~4 mol dm~3, (=),
2
with the spectrum of HO ~/O ~ at pH 4.0, (L) (see text)
2
2
CHO
~OOCH(OH)SO ~ ] HO ~ ] =
(6)
3
2
SO ~
3
A possible alternative mode of decay of ~OOCH(OH)SO ~
3
is that SO 2~ acts as a leaving group to yield HSO ~ and
3
3
~OOCHO, reaction (7).
~OOCH(OH)SO ~ ] ~OOCHO ] HSO ~
(7)
3
3
To explore further the decay of ~OOCH(OH)SO , it was
deemed appropriate to study the release of H` using conduc-
Fig. 3 (a) Conductivity signal developing over 1 ms in oxygenated
solution at pH 4.0. The superimposed line is calculated using FAC-
SIMILE (see text). (b) Conductivity signal developing over 10 s in
oxygenated solution at pH 4.0.
3
tivity detection. As conductivity measurements can not readily
be made at high salt concentrations, the concentration of
HMS in these experiments was reduced to 12.5 mmol dm~3.
Increases in conductivity were observed on two timescales, see
Fig. 3(a) and (b). Fig. 3(a) indicates a biphasic growth of con-
ductivity, the fast initial rise of which corresponds approx-
software package23 to a scheme involving reactions (6), (7) and
(8) and the dissociation of HO ~ and formic acid in equilibria
2
(5) and (9). The Ðtting procedure assumed that k ] k \ 2.4
imately with the rate of decay of ~OOCH(OH)SO ~
6
7
3
] 104 s~1, i.e. the value obtained from the optical measure-
determined optically. The extent and kinetics of the conduc-
ments for the decay of ~OOCH(OH)SO ~. The additional
tivity changes shown in Fig. 3(a) are compatible with the
3
parameters used in the curve-Ðtting are listed in Tables 1 and
decay of ~OOCH(OH)SO ~ proceeding by a split path involv-
3
2 and the process yielded k \ 1.7 ] 104 s~1, k \ 7.0 ] 103
ing reactions (6) and (7) if it is also assumed that the product
6
7
s~1 and k \ 2.5 ] 103 s~1. The ratio of k to k obtained is
of reaction (6), i.e. ~OOCHO, can react to release further H`.
We propose that the process involved is the hydrolysis of
~OOCHO to yield the hydrated peroxyformyl radical, which is
8
6 7
somewhat dependent on the value of j
reducing j ~ by 10% resulted in k increasing by about 2.5%.
used. For example,
~
O2
O2
6
known to decompose in \1 ls to formic acid and HO ~, reac-
HCOOH H H` ] HCOO~
(9)
2
tion (8).22 This reaction also provides one route to the forma-
The low value of E and negative entropy of activation
tion of the formate, which is observed as a product (see
below).
a
[*S” \ [(68.2 ^ 4.2) J K~1 mol~1] obtained from the tem-
perature dependence of the decay of ~O CH(OH)SO ~ is in
2
3
~OOCHO ] H O ] ~OOCH(OH) ]
keeping with the heterolytic nature of reactions (6) and (7).
The slow growth of conductivity over a period of seconds
illustrated in Fig. 3(b) was not studied in detail, but corre-
sponds with the release of H` with G(H`) B 6.7 ] 10~7 mol
J~1, probably by several concurrent processes (see below).
2
2
HCOOH ] HO ~ (8)
2
The line superimposed on the trace in Fig. 3(a) results from
Ðtting the experimental trace using the FACSIMILE kinetics
Table 1 Reactions and rate constants used in the FACSIMILE programme for Ðtting the “fastÏ changes in conductivity
reaction
k
reference
~CH(OH)SO ~ ] O ] ~O CH(OH)SO ~
2.6 ] 109 dm3 mol~1 s~1
1.7 ] 104 s~1
this work
this worka
this worka
3
2
2
3
~O CH(OH)SO ~ ] HO ~ ] CHOSO ~
2
3
2
3
~O CH(OH)SO ~ ] ~OOCHO ] HSO ~
7.0 ] 103 s~1
2
3
3
~O CHO ] H O ] HO ~ ] HCOOH
2.5 ] 103 s~1
this worka
2
2
2
HO ~ H O ~ ] H`
1.0 ] 106 s~1, 3.6 ] 1010 dm3 mol~1 s~1
9.0 ] 106 s~1, 3.6 ] 1010 dm3 mol~1 s~1
20, corrected for high ionic strength
2
2
HCOOH H HCOO~ ] H`
K
corrected for high ionic strength
diss
a Result of FACSIMILE Ðt.
J. Chem. Soc., Faraday T rans., 1997, V ol. 93
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Table 2 Other parameters used in the FACSIMILE programme for
of HO ~ with the HSO ~, which is in equilibrium with HMS,
2
3
Ðtting the “fastÏ changes in conductivity
and (ii) the ~OH radical-initiated chain oxidation of HSO ~.11
3
The Ðrst of these is unable to account for the full yield of
parameter
value
reference
SO 2~ since, on the mechanism given above for the decay of
4
~OOCH(OH)SO ~, the yield of sulfate would be given by
G(~OOCH(OH)SO ~)
5.6 ] 10~7 mol J~1
6.0 ] 10~8 mol J~1
0.0335 S m2 mol~1
0.0065 S m2 mol~1
0.005 45 S m2 mol~1
0.0040 S m2 mol~1
see text
see text
24
estimated
24
3
3
G(H~)
G(SO 2~) \ G(H O )
j
`
4
2 2
H
j
~
O2
\ 1G(HO ~) ] G
j
2
2
H2O2
HCOO~
j
24a
~
\ 1[G(~OH) ] G ] G
]
HSO3
2
H
H2O2
a Assumed the same as HCO ~ and corrected to 20 ¡C.
B 4.0 ] 10~7 mol J~1
3
where G and G
from the radiolysis of water. The yield of SO 2~ by (ii) is diffi-
cult to predict, but it appears, on the basis of the small frac-
are the primary yields of H~ and H O
H
H2O2
2
2
4
Products of c-radiolysis at pH 4.0
tion of ~OH that reacts with HSO ~, that initial chain lengths
3
are approximately 50 and 70, respectively, at the two dose
The yields of formate and sulfate formed on c-irradiation of
rates. The non-linear formation of SO 2~ is probably due to
solutions with [HMS] \ 0.01 mol dm~3, [N O] \ 0.0125
4
2
mol dm~3 and [O ] \ 6.75 ] 10~4 mol dm~3 were studied
2
the formate product interfering with the chain reaction as a
result of it scavenging radicals involved in the chain propaga-
for doses up to 1500 Gy. Experiments were carried out at two
tion, e.g. SO ~~.
dose rates, namely, 16.3 and 2.3 Gy min~1. The yields of
formate were proportional to the absorbed dose and were
independent of dose rate, see Fig. 4. The gradient of the graph
indicates that G(formate) \ (7.1 ^ 0.2) ] 10~7 mol J~1, which
is somewhat higher than G(~OH) at this concentration of
HMS (5.7 ] 10~7 mol J~1). This yield of formate also implies
4
Reaction of ~CH(O—)SO — with oxygen
3
Absorption spectrum of the product. The reaction of
~CH(O~)SO ~ with oxygen was investigated at pH 7.0 using
3
solutions containing 0.1 mol dm~3 HMS that were bu†ered
that CHOSO ~, the product of reaction (6), must also lead
3
with 0.1 mol dm~3 phosphate bu†er, thus causing equilibrium
(3) to be established in 120 ns.13 Following the pulse, an
absorption grew over 6 ls for wavelengths between 240 and
300 nm and then decayed, only slowly, over millisecond time-
scales. The spectrum measured after completion of the growth
is shown in Fig. 5. This spectrum is not due to
ultimately to the formation of formic acid. Hydrolysis of
CHOSO ~ to HCOOH and HSO ~, reaction (10), is a plaus-
3
3
ible pathway for this and this process is a possible contributor
to the release of H` illustrated in Fig. 3(b).
CHO
CH(OH)
2
=
] H O ] =
] HSO ~ ] HCOOH (10)
~OOCH(OH)SO ~ (cf. Fig. 1, B), but corresponds very closely
2
3
SO ~
SO ~
3
3
3
with that of O ~~,19 which is shown in the Ðgure for compari-
2
son. This spectrum is calculated assuming a yield given by
Unlike the formate, the yields of sulfate depend on dose rate
and there is some indication of departures from linearity at
the upper extremity of the dose range studied (see Fig. 4). The
G(O ~) \ G[~CH(O~)SO ~] ] G(HO ~)
2
3
2
slopes of the yield vs. dose plots give G(SO 2~) \ 2.0 ] 10~6
\ G(~OH) ] G(H~)
4
mol J~1 at a dose rate of 16.3 Gy min~1, and 2.7 ] 10~6 mol
where G(~OH) \ 6.0 ] 10~7 mol J~1 at the concentration of
HMS and G(H~) \ 6.0 ] 10~8 mol J~1. This suggests that
~CH(O~)SO ~ reacts directly with O to yield O ~ [reaction
J~1 at a dose rate of 2.3 Gy min~1. Both yields are consider-
ably in excess of G(~OH), but as discussed above there is no
evidence to suggest that SO 2~ is generated by the decompo-
3
2
2
4
(11)] and that if a peroxy radical is formed as an intermediate
its lifetime is very short.
sition of ~OOCH(OH)SO ~. Plausible mechanisms for the for-
3
mation of SO 2~ are: (i) reaction of the hydrogen peroxide
formed by the radiolysis of water and the disproportionation
4
~CH(O~)SO ~ ] O ] O ~~ ] CHOSO ~
(11)
3
2
2
3
Further evidence that the long-lived product is O ~~ was
sought by studying the decay kinetics of the product. Since the
2
Fig. 4 E†ect of dose and dose rate on the yields of formate and
sulfate ions formed on c-radiolysis of oxygenated HMS solutions at
pH 4.0. Yields of formate ion and sulfate at dose rates of 2.3 Gy
min~1, (=) and (K), and at 16 Gy min~1, (…) and (>), respectively;
Fig. 5 Comparison of the spectrum obtained on pulse radiolysis of
oxygenated solutions at pH 7.0 (|) with that of O ~~ (the solid line is
2
[O ] \ 6.75 ] 10~4 mol dm~3.
calculated as described in the text); [O ] \ 6.75 ] 10~4 mol dm~3
2
2
3644
J. Chem. Soc., Faraday T rans., 1997, V ol. 93

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decay of O ~~ is known to be sensitive to the presence of
who made and developed the ac conductivity detection
system.
2
transition metal ions, the decay of absorption was monitored
at 250 nm in the presence of 100 lmol dm~3 EDTA to
complex any metal ions. The decay was second-order with 2k/
e \ 9.7 ] 102 cm s~1 which, taking e (O ~~) \ 1890 dm3
References
1
J. W. Munger, C. Tiller and M. R. Ho†man, Science, 1990, 86,
545.
250
2
mol~1 cm~1,19 yields 2k \ 1.8 ] 106 dm3 mol~1 s~1. This
value is close to, but slightly higher than, the literature value20
of 1.4 ] 106 dm3 mol~1 s~1 for the disproportionation of
2
E. G. Chapman, C. J. Barinaga, H. R. Udseth and R. D. Smith,
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3
4
P. Warneck, J. Atmos. Chem., 1989, 8, 99.
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2
G. L. Kok, S. N. Gitlen and A. L. Lazrus, J. Geophys. Res., 1986,
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k
B 33k , and the slightly higher value of the rate constant
13
12
observed here is probably due to the inÑuence of the second-
5
6
L. W. Richards, J. A. Anderson, D. L. Blumenthal, J. A. McDon-
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4863.
ary salt e†ect on K (HO ~) owing to the high ionic strength
diss
2
of the solutions.
HO ~ ] HO ~ ÈÈÈ H O ] O
(12)
(13)
7
8
9
W. J. McElroy, Atmos. Environ., 1986, 20, 323.
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2
2
2
2
`
H
HO ~ ] O ~~ ÈÈÈ H O ] O
2
2
2 2
2
10 G. V. Buxton, S. McGowan, G. A. Salmon, J. E. Williams and
N. D. Wood, Atmos. Environ., 1996, 30, 2483.
Rate of reaction (11). The rate of reaction (11) was studied
by pulse radiolysis with 15 Gy, 200 ns pulses of solutions con-
taining 0.1 mol dm~3 HMS, 0.1 mol dm~3 phosphate bu†er,
oxygen in the range (2È11) ] 10~4 mol dm~3 and nitrous
oxide in the range (4.6È20) mmol dm~3. The formation of
11 U. Deister, P. Warneck and C. Wurzinger, Ber. Bunsen-Ges. Phys.
Chem., 1990, 94, 894.
12 U. Deister, R. Neeb, G. Helas and P. Warneck, J. Phys. Chem.,
1986, 90, 3213.
13 S. Barlow, G. V. Buxton, S. A. Murray and G. A. Salmon, J.
Chem. Soc., Faraday T rans., 1997, 93, 3637.
O ~~, monitored at 250 nm, obeyed Ðrst-order kinetics and
2
14 J. W. J. Spinks and R. J. Woods, in An Introduction to Radiation
Chemistry, Wiley, New York, London and Sydney, 3rd edn.,
1990.
the observed rate constant, k , was proportional to [O ]
obs
2
with slope k \ (1.6 ^ 0.1) ] 109 dm3 mol~1 s~1.
11
15 J. Grimshaw, J. R. Langan and G. A. Salmon, J. Chem. Soc.,
Faraday T rans., 1994, 90, 75.
Conclusions
16 E. Janata, J. Lilie and M. Martin, Radiat. Phys. Chem., 1994, 43,
353.
At pH 4.0, ~OOCH(OH)SO ~ dissociates by a split pathway,
3
17 D. Veltwisch, E. Janata and K-D. Asmus, J. Chem. Soc., Perkin
reactions (6) and (7), but the Ðnal products are in each case
T rans. 2, 1980, 2, 146.
formic acid and HO ~, which in turn disproportionates to
2
18 E. Bothe, M. N. Schuchmann, D. Schulte-Frohlinde and C. von
Sonntag, Photochem. Photobiol., 1978, 28, 639.
19 A. J. Elliot and G. V. Buxton, J. Chem. Soc., Faraday T rans.,
1992, 88, 2465.
H O and O , but we have found no evidence for a chain
2
2
2
production of formate as suggested by Deister and Warneck.8
The hydrogen peroxide is expected to form sulfuric acid by its
reaction with the HSO ~, which is in equilibrium with the
20 B. H. J. Bielski, D. E. Cabelli, R. L. Arundi and A. B. Ross, J.
Phys. Chem. Ref. Data, 1985, 14, 1041.
21 R. P. Bell, Acid-Base Catalysis, Clarendon Press, Oxford, 1941.
22 E. Bothe and D. Schulte-Frohlinde, Z. Naturforsch. B: Chem.
Sci., 1980, 35, 1035.
23 A. R. Curtis and W. P. Sweetenham, FACSIMIL E/
CHECKMAT Users Manual, UKAEA, AERE, 1987, R12805.
24 R. A. Robinson and R. H. Stokes, Electrolyte Solutions, Butter-
worths, London, 2nd edn., 1959.
3
HMS. Sulfate is also generated by an additional path, which is
most likely the free-radical chain oxidation of HSO ~.
3
At pH 7, the ~CH(O~)SO ~ radical reduces O directly to
3
2
O ~ with no evidence for the intermediate formation of an
2
RO ~ radical.
2
We acknowledge Ðnancial support from the Commission of
the European Union through research contracts STEP-0005-
C(MB) and EV5V-CT93-0317. We also thank Mr C. Kilner,
Paper 7/03266H; Received 12th May, 1997
J. Chem. Soc., Faraday T rans., 1997, V ol. 93
3645