Electrolyte effects on aqueous atmospheric oxidation of sulphur dioxide by hydrogen peroxide

Atmospheric Environment VoL 27A, No. 2, pp. 129-137, 1993. 0004 6981/93 $6.00+0.00 Printed in Great Britain, © 1993 Pergamon Press Ltd

ELECTROLYTE EFFECTS ON AQUEOUS ATMOSPHERIC OXIDATION OF SULPHUR DIOXIDE BY HYDROGEN

PEROXIDE

J. LAGRANGE, C. PALLARES, G. WENGER and P. LAGRANGE URA 405 au CNRS, Laboratoire de Cin~tique et Analyse, EHICS, 1 rue Blaise Pascal, F-67000 Strasbourg,

France

(First received 2 June 1992 and in final form 23 September 1992)

Abstract--The rate of oxidation of sulphur dioxide by hydrogen peroxide in the presence of various supporting electrolytes has been determined by the stopped flow method. In a sodium perchlorate medium (chosen as reference electrolyte) first-order kinetics were observed with respect to HSO~, H202 and H +. The influence of the ionic strength was investigated and the third-order rate constant was calculated at infinite dilution. The rate constants of the reaction are shown to be far higher when chloride or ammonium ions are added to the solution.

The effect of temperature and of traces of metal ions (Mn 2 +, Fe 2 +, Fe 3 +, Cu 2 +, Cr 3 +) was studied. The oxidation reaction is virtually insensitive to the effects of Mn 2 +, Fe 3 +, Cu 2 + and Cr 3 +. Catalytic activity is observed with Fe 2 +.

Key word index: Hydrogen peroxide, cloud, acid, oxidation, kinetics,

INTRODUCTION

The occurrence of sulphuric acid, present in the atmosphere and in precipitation, results mainly from the oxidation of SO2 which is emitted from either natural or man-made sources in the troposphere. Oxidation mechanisms include not only homogeneous gas-phase reactions, but also multiphase processes occurring in clouds. An important portion of sulphate found in precipitation is due to aqueous-phase oxidation of SO2 dissolved in cloud or rain water. The main oxidative processes of dissolved SO2 involve atmospheric oxygen, ozone or hydrogen peroxide. Recent laboratory studies have shown that H202 is probably the key atmospheric oxidant of SO2 in precipitation, if the pH is less than 6 (Mader, 1958; Hoffman and Edwards, 1975; Penkett et al., 1979; Martin and Dam- schen, 1981; Lee et al., 1986; Lind et al., 1987; Saxena and Seigneur, 1987; Meagher et al., 1990; Radojevic et al., 1990; Fung et al., 1991). However, most of these studies involved dilute solutions.

In this paper we present some recent data on the kinetics of oxidation of dissolved SO2 by H202, in solutions containing various ions (CI-, SO 2-, NH~') at high ionic strengths and metallic traces (Mn 2+, Fe 2+, Fe 3+, Cu 2+, Cr 3+) which are relevant to the conditions occurring in the atmosphere. The discussion of these parameters is essential when considering results published in the literature (Lind et al., 1987; Chandler et al., 1988). These results show that field measurements lead to a rate of oxidation of SO2 by H202, larger than the value obtained by laboratory measurements. These results indicated the influence of trace components on SO2 oxidation: me-

tallic traces and the main anions or cations found in water droplets. High concentrations of electrolytes may be observed under atmospheric conditions, in particular in an aerosol of deliquescent particles.

In our kinetic study, a series of experimental measurements were carried out at a given constant ionic strength with a given supporting electrol3,te. For each experiment the concentration of the supporting electrolyte is sufficient to ensure that the activity factors of the reactants remain at constant values. The rate laws are expressed, as usual, in function of the concentration of the reactants. Thus, we had to determine the concentration of the reactants for each kinetic test solution. In the case of the H + reactant, our technique allows the measurement of the concentration of H ÷ instead of the activity.

EXPERIMENTAL

Reagents

All chemicals were of analytical reagent grade. Sodium perchlorate, sodium sulphate, sodium chloride, sodium acetate, sodium monohydrogenophosphate and perchloric acid were obtained from Merck. Sodium sulphite, hydrogen peroxide, sodium hydroxide, hydrochloric acid and sulphuric acid were purchased from Prolabo. Analytical grade F e ( C 1 O a ) 2 , Fe(CIO4) a (Aldrich), Cu(CIO4) 2 (Fluka), (CHaCO2)TCr3(OH)2 (Aldrich), and Mn(CHaCO2) 2 (Al- drich) were used without further purification.

The analytical concentration of sulphite was determined by iodometric titration. The concentration of hydrogen peroxide was determined by potentiometry (redox reaction with Ce 4+) using a standardized solution of ferrous ions in sulphuric medium (Chariot, 1974).

129

t30 J. LAGRANGE et al.

Oxygen was displaced by bubbling argon through the solution before analytical and kinetic measurements were taken.

H + concentration measurements

[H +] measurements were carried out with a Tacussel TB/HA glass electrode and an Ag/AgCI/0.1 molt ' -1 CI- reference electrodo at 25+0.1°C. The ionometer was a Tacussel Ionoprocesseur II. The glass electrode was calib- rated against a strong acid solution (10- 2 tool ~'- 1 HCIO 4 in NaCIO4 or in NH4CIO 4, or 10 -z mol E- 1 HCI in NaCI), for each ionic strength and supporting electrolyte. For these standard solutions, we took - log [H +] = 2.000. In Na2SO 4, the standardization of the glass electrode is performed in basic medium. In Na2SO 4 (1=0.6, as example), a solution 0.01 tool E- 1 M NaOH + 0.1967 mol ~'- 1 Na2SO4 was used as a standard. Since in this medium the apparent constant K , = [ H +] [ O H - ] and pK,, = 13.07 at 25°C (Fischer, 1967), we took - l o g [ H + ] = 11.07 for this standard.

Thus in our work, in a given ionic medium, the logarithm of the proton concentration was measured. In the literature, the definition of the pH scale concerning the kinetics of oxidation of sulphite by H202 is never dearly indicated. Only Drexler et al. (1991) pointed out that their pH measurements lead to the measurement of the proton activity.

The solutions were generally buffered by addition of sodium acetate (0.04 mol g- 1) or sodium phosphate (0.04 mol E- 1) and [H +] was adjusted before kinetic measurements by the addition of sodium hydroxide or perchloric (or hydrochloric or sulphuric) acid. The H + concentrations of the mixtures were measured after each kinetic measurement and these measurements have shown the sufficient buffer capa- city of all studied kinetic mixtures.

Kinetic measurements

The stopped-flow spectrophotometer, type Durrum Gib- son equipped with a transcient recorder Datalab DL 905, was interfaced to an Apple II microcomputer. On-line ac- quisition and treatment of data have been previously described (Lagrang¢ and Lagrange, 1984). In all kinetic studies, the H + concentration was kept constant by the addition of a buffer (acetate or phosphate: - l o g [H +] in range 3-6) or of a strong acid ( - l og [H+]=3) . In all kinetic studies, solutions of H20 2 (10-4-5x 10 -4 mol g-l), in a given electrolyte medium and at fixed H + concentration, were mixed with solutions of sulphite (10-a-2 x 10 -2 mol t ' - 1) in the same medium and at the same H + concentration. Sul- phite ions were present in sufficient excess to ensure that the observed step of the sulphite oxidation was always pseudo first-order with respect to hydrogen peroxide. The measurements were carded out at 25_0.I°C and at 245 nm. Each value of the pseudo first-order rate constant (koch) was the average of at least three determinations.

RESULTS AND DISCUSSION

Studies in NaC104 medium

Determinat ion o f the rate law in NaCIO4 medium. Experiments were performed in the range 3 < - l o g [ H +] < 5.5 where the p redominan t species of S(IV) is HSO~. Unde r these condi t ions the kinetic law proposed in the l i terature (Mader , 1958; Hoffmann and Edwards, 1975; Penket t et al., 1979; Mar t in and Dam- schen, 1981; Lee et al., 1986; Lind et al., 1987) for the oxidat ion of dissolved SO2 by H 2 0 2 is given by:

v = k [ H 2 0 2 ] [ H S O ~ ] f ( [ H + ]), (1)

where f ( ) is a function.

Usually, first-order kinetics are found with respect to hydrogen sulphite and hydrogen peroxide. How- ever, the function f ( [ H + ] ) is not perfectly defined. It can be either I-H +] or a more complex function of [H + ]. The different results reported are probably due to different experimental conditions: kinetic runs performed with various buffers and with or without a suppor t ing electrolyte; the term [H +] used for the kinetic analysis is either the pro ton concent ra t ion or the pro ton activity.

In our study, drastic experimental condi t ions are imposed to avoid some effects due to the var ia t ion of the ionic envi ronment , which are poorly discussed in the literature. To determine the parameters of the kinetic law, the ionic s t rength and the ionic environment were mainta ined constant with a suppor t ing electrolyte at high excess concent ra t ion (25 times in excess) vs buffer concentrat ion. The concentra t ions of the reactants (H202 and sulphite) are always 50 times lower than the suppor t ing electrolyte concentrat ion. The chosen suppor t ing electrolyte is NaCIOa at 1 mol f-- 1. This suppor t ing electrolyte is known to give very low interactions with ions and molecules.

For each of the kinetic runs, the plots of In I A, - A ~ I vs time were always linear for the selected time range. The parameter A, is the absorbance of the system at reaction time t, and A~, the absorbance at t~. These results confirm first-order kinetics with reference to H202 and lead to the values of kobs, the pseudo first-order rate constant .

Fo r a given cons tant - l o g [H + ] = 4.8, experiments are performed with acetate and phospha te buffer and with various concentra t ions of sulphite (between 2 × I0 - 3 and 1.5 x 10- 2 mol f - 1). In these two cases, the pseudo first-order rate cons tant varies linearly with [ H S O 3 ] (Fig. 1) with different slopes for acetate and phospha te buffer. These results show first-order kinetics with respect to HSO~- and a buffer effect on the rate constant .

16-

12,

kob s 8' (s" 1 )

4 ̧

o~ o

PHOSiHATE ~ AC~/TATE

4 8 12 Sulphite concentration (mmol I "1)

Fig. 1. Pseudo first-order rate constant (kob~) as a function of [HSO3-] at - log [H +] =4.8 in acetate or phosphate buffer; ionic s trength=l tool

f 1 (NaCIO4): 25°C

Atmospheric oxidation of sulphur dioxide 131

To determine the function f ( [H+]) (Equation 1), kinetic runs at constant [H +] (3 < - l o g [H +] < 5.5) with the concentration of sulphite quasi constant (between 3x10 -3 and 5x10 -3 moi f - l : [HSOf]a.) were performed at constant ionic strength in NaCIO4 medium (1 mol d-1). At - l o g [H +] > 3.5, additions of a buffer (acetate or phosphate) were always em- ployed to control the H + concentration. At - l o g [H+]=3 , kinetic runs were performed with buffers (acetate or phosphate) and without buffer. For each kinetic run, the pseudo first-order rate constant was measured, and a corrected pseudo first-order rate constant was calculated for a constant [HSO3] = 4 x 10- 3 mol f - 1 (first-order kinetics with reference to HSO3):

kcor = (ko~ x (4 x 10- 3))/[HSO3 ],, .

The logarithm of this corrected pseudo first-order rate constant in function of - l o g [H + ] is given in Fig. 2. The two parallel straight lines ( s lope=-1) obtained with an acetate or phosphate buffer indicate a first-order reaction with respect to H +, and a buffer effect on the value of the rate constant. All these results lead to the following kinetic law:

v=k [H202] [HSO~] [H+] . (2)

In our experiments where the buffer concentration is low but sufficient to keep [H +] constant ([buffer] ~< 0.04 mol .f- 1), the third-order rate constant, k, is independent of the concentration of the buffer and of the H ÷ ions. But k varies with the nature of the buffer. All kinetic data are summarized in Table 1.

The k value is practically the same with acetate buffer as without buffer. Thus, acetate buffer was chosen to study the other parameters: influences of the nature and the concentration of the supporting electrolyte, of the temperature and of the metallic traces. This kinetic law is checked in the other studied ionic media (NaC1, Na2SO4 and NH4C104).

There is a lack of information on the mechanism of the oxidation of sulphite by H202, only a few studies have proposed a mechanism. Hoffmann and Edwards (1975) and recently Drexler et al. (1991) suggest the formation of the peroxymonosulphurous acid intermediate. This intermediate is subject to acid catalysis (rate determining) with protons donated either from H30 + or from the buffer.

Fast pre-equilibrium:

- O

HSO~ + H 2 0 2 , ' S -4 ) -OH+H20 . /

O

Specific acid catalysis (rate-determining step):

-O \ k,

S-O4DH + H ÷ ,SO~- + 2 H ÷. /

O

Generalized acid catalysis (rate-determining step) due to the presence of a buffer solution (HA/A-):

- O k "

S-O-OH + AH ,SO~- + 2 H + + A - . /

O

Hence the rate expression is written as follows:

v = [H202 ] [HSO~ ] (Kk' [H + ] + Kk" [HA])

with the third-order rate constant k=Kk '+Kk" [HA]/[H+].

The third-order rate constant, k, found by this mechanism appears as a function of the H +concen-

2-

log kco r 1-

- 1 i i

zs 3'.s ,'.s I s - I o 0 1 8 + ]

Fig. 2. Logarithm of the corrected pseudo first-order rate constant (k~or) as a function of - log [H +] in acetate or phosphate buffer; [HSO~]=4x 10 -3 tool

: - 1; ionic strength = 1 mol ~'- 1 (NaCIO4); 25°C.

Table 1. Influence of the buffer on the third-order rate constant in NaCIO4 at 1 mol #- 1 at 25°C

Buffer Studied concentration Studied [HSO~ ] k (mol d-l) --logI-H +] (mol :-1) (s-1 mol-2 :z)

Without 3.30 0.0025 (3.07+0.10) × 107

CH3COOH/CH3COO- 3.00-5.50 0.003-0.015 (2.80+0.10) x 107 0.02-0.04

Phosphate 3.00-5.50 0.003-0.015 (7.10-1-0.30) x 10 ~ 0.02-0.04

132 J. LAGRANGE et al.

tration and of the concentration of the HA acid form of the buffer. But our experimental data lead to a value of k independent of H + and buffer concentrations. Generalized acid catalysis does not seem to be taken into account in the proposed mechanism of Hoffmann and Edwards (1975). As the value of the rate constant is practically the same in acetate medium and without buffer, the effect of this buffer may be excluded. Then with acetate buffer, k = Kk'.

In phosphate medium where the value of the third- order rate constant (independent of H ÷ and buffer concentrations) is approximately three times higher than that with acetate buffer, generalized acid catalysis might also be neglected and an oxygen atom transfer catalyst giving two pre-equilibria may be proposed:

K' phosphate + H202 ~ peroxyphosphate

1.5-

k 1 - ~ (108M-2~ 1)

0.5,

0 01~5 015 0.~5 ' Ionic strength (tool I "1)

Fig. 3. Influence of the ionic strength on the third-order rate constant (k) in NaCIO4 medium. * Experimental

data; the,curve is calculated using Equation (3).

O K" \ peroxyphosphate+HSO~ ~ S-O-OH+phosphate.

/ O

These two pre-equilibria are followed by specific acid catalysis and the third-order rate constant may be expressed as:

k=k 'K 'K" .

These mechanisms proposed are supported by experimental data obtained at constant ionic environment (the concentration of the buffer 25 times lower than the concentration of the supporting electrolyte) and constant ionic strength.

Some authors, in particular Hoffmann and Edwards (1975) and Drexler et al. (1991), found a generalized acid catalysis due to the presence of the acid form of the buffer. But these authors work with a buffer concentration 20 times higher than in our study. Under these conditions, if the ionic strength is kept constant, the ionic environment varies with the concentration of the buffer (acetate ion may represent as much as 80% of the ionic strength). Our data show, clearly, an effect of ionic environment at constant ionic strength in NaCI-NaCIO4 medium as described in the following section.

Furthermore, our experimental results are always given in terms of the concentration of each reactant as usual in the kinetic studies. On the other hand, Drexler et al. (1991) give a rate law dependent on concentration of the reactants (except H ÷) and the activity of H +. Our work does not indicate a mechanism with generalized acid catalysis. This effect, apparent in some literature data, may instead be due to the variation of the ionic environment.

Effects of the ionic strength. The third-order rate constant of Equation (2) was determined for several concentrations of NaCIO4, and the results are shown in Fig. 3. The variation of the third-order rate constant with ionic strength exhibits a maximum for

a concentration of NaCIO4 of about 0.3 mol E- i. This variation of the rate constant suggests an effect of the ionic strength. The third-order rate constant (k = k'K) appears in the mechanism as a product of an equilibrium constant and a rate constant. In the activated complex theory, the rate constant may be expressed as a function of the equilibrium constant of the activated complex. These equilibrium constants may be expressed as a function of the activity factors of the reactants. The equation of Debye-H/ickel (with the hypothesis of Guggenheim), which gives a good rep- resentation of the variation of the activity factor of an ion, i, in a concentrated electrolyte solution containing more than one electrolyte species (ions j), may be introduced in the expression of the equilibrium constants. The rate constant can be expressed as:

log (k/ko)=A'xSI/(1 + x / l ) + ~ F, jC,, (3)

where A is the function of the absolute temperature, of the dielectric constant, and of the charge of the ions i; F 0 is the adjustable parameter characteristic of the ions i and j; Cj is the concentration of ion j of the supporting electrolyte; I is the ionic strength; and k0 is the rate constant at infinite dilution.

The parameters of this equation can be calculated with the program Letagrop Modfunction:

ko=0.672x 107 s lmol -2f2

Fi; = - 2.05.

Usually the variation of the mean activity factor of various electrolytes passes through an extremum in the range of the ionic strength I=0.1 - I mol ~ - 1 A maximum value of log k is observed in the same range of concentration. When I > 0.6 mol f - l the F, FoC j term becomes predominant: effectively log k decreases quasi-linearly with the ionic strength.

Effects of the ionic supporting electrolyte

NaC1 medium. The third-order rate constant, k, as a function of the ionic strength in NaCI is shown in


Fig. 4. The k value is nearly constant as a function of the ionic strength. But if we take the ko value obtained at infinite dilution in NaCIO4 medium, the parameters of Equation (3) may be calculated in NaCI medium.

8.5-

log k

7.5,

6.~

f I ' i i

0.25 01.5 1 0.75 Ionic strength (mol I "1)

Fig. 4. Influence of the ionic strength on the third- order rate constant (k) in NaCI medium. + Experi- mental data; the curve is calculated using Equation (3).

A catalytic effect of CI- ions on the kinetics of oxidation of dissolved SO2 by 02 has been discussed by Clarke and Radojevic (1983). The quasi-constant value of k is probably due to two opposing factors:

• variation of the ionic strength leading to a decrease of the rate constant when the XFuC ~ term becomes predominant,

• a catalytic effect of the Cl- ion leading to a linear increase of the rate constant in function of Cl- concentration. The EF U ferm (Equation 3) in Nael medium is found to be:

F u = - 1.20.

NH4CIO, medium. Values of the third-order rate constant obtained in NH,C104 medium are given in Table 2. In NH4CIO4 medium, similar effects to those obtained in NaC1 medium are observed with approxi- matively the same value of the rate constant when the ionic strength is varied.

Na2SO, medium. Values of the third-order rate constant obtained in Na2SO4 medium are given in Table 3.

Table 2. Influence of the ionic strength: experimental results for NH,C10, medium (acetate buffer) at 25°C

I [HSO3] x 103 /Cob ~ k x 10 -~ (tool d -1) - log[H + ] (mol f - t ) (s-t) (s-t moi-2 d2)

1.0 4.50

0.1 4.50

31.9±0.6 9.65 32.2±0.6

29.5±0.5

9.8±~2 5.~ 9.~ 11.3±0.3

10.5±0.2

25.7±0.5 9.07 25.5±~6

27.1±0.6

7.49±0.15 5.~ 9.07 7.55±~16

7.67±0.16

10.6±0.5

8.70±0.40

Table 3. Influence of the ionic strength: experimental results for Na2SO 4 medium (acetate buffer) at 25°C

1 [HSO~] x 103 ko~ k x 10 -7 (moi d -l) --Iog[H + ] (mol g-l) (s-l) (s-i mol-2 d2)

1,68 ± 0.04 0.4 5.00 5.45 1.78+0.04 3.14+0.19

1.73 ± 0.04 1.75±0.03

1.17+0.02 5.00 5.42 l.l 7 ± 0.02

1.17±0.03 1.18±0.03

1.07 ± 0.02 !.0 4.98 5.42 1.08 ± 0.03

1.09 ± 0.03 1.11 ±0.02

0.6 2.17±0.13

1.90+0.10


Results for ionic strength less than 0.4 M are not reported in Table 3. At low ionic strength ( < 0.4 M), the concentration of acetate buffer vs sulphate concentration is not negligible and when the ionic strength is kept constant, the ionic environment varies.

In NazSO4 medium, similar effects to those obtained in NaC104 medium are observed. At high ionic strength, in acetate buffer, a decrease of the rate constant is observed when the ionic strength increases.

NaCIO4-NaCI medium at constant ionic strength. The variation of the third-order rate constant at constant ionic strength (mixture of NaC1 and NaCtO4) as a function of [C1-] is shown in Fig. 5. The curve exhibits a maximum for a concentration of 0.4 mol { - ' in NaCI.

At constant ionic strength (1 mol E-~) in a NaC1- NaCIO4 mixture, Equation (3) should lead to a linear variation of log k vs CI- concentration. This effect is not observed. The nature of the supporting electrolyte is not sufficient to explain the observed variation (Fig. 5) of the rate constant experimentally obtained. A catalytic effect might also be taken into account. The chloride ion should give an intermediate species in fast pre-equilibria. When the concentration of CI- increases, intermediates with more than one chloride ion (probably less reactive) should also be taken into account.

An empirical law may be proposed:

kNaCIO, + NaCI = kNaCIO, + kNaCl, (4)

where kN,c,o,+Nac~ is the third-order rate constant measured in a mixture of NaC1 (C, mol E -1) and NaCIO4 (C2 mol E-I) ; C~ + C 2 = c o n s t a n t = 1 mol f - 1; kNacl is the third-order rate constant measured in NaC1 (C1 mol ( - 1); and ks ,oo, is the third-order rate

constant measured in NaC104 (C2 mol {-1). Drexler et al. (1991) have studied the influence of

the ionic medium on the kinetics of oxidation of sulphite by H202. These authors have not observed

a kinetic anion effect for ClOg, C1- and SO 2-. But they introduced in their kinetic law the proton activity instead of proton concentration. This difference may explain the different conclusions on the effect of the ionic strength. Our definition of rate laws (expressed in function of the reactant concentrations) agrees better with the definition usually used in the literature to examine the influence of the ionic medium in kinetic studies.

Temperature effects

The third-order rate constant was also determined at various temperatures (13, 18, 25 and 35°C) at 4.7 < - l o g [H +] < 5.1, in NaCIO4 medium with acetate or phosphate buffer, in NaCI and Na2SO4 medium with acetate buffer (Tables 4"-6). The Arrhenius activation energy (Ea), the entropy (AS:~) and the enthalpy (AH~) of activation were calculated. These

log k

8.5.

7.5:

EXPERIMENTAL DATA ~ _ _ _ADI/c~/VE LAW

LAW

o.~'s o'.s o.~'s Chloride ion concentration (mol I "1)

Fig. 5. Influence of [CI- ] on the third-order rate constant (k) at constant ionic strength (1 tool (-1) in a mixture of NaCI-NaC104. ~- Experimental data; ( . . . . ) experimental curve; ( . . . . ) calculated curve with the empirical law (Equation 4); ( . . . . . ) calculated curve with the Debye-Hiickel equation (Equa-

tion 3).

Table 4. Temperature effects in NaCIO 4 medium. The third-order rate constant (k) is a mean value of at least three experimental values of kob s. The c o n s t a n t kfi t is the third-order rate

constant calculated from the Arrhenius law

Medium I Temp. [HSO~] x 103 k x 10 -7 kri , x 10-7

(buffer) (°C) - l o g [ H +] (tool (-1) (s-t mol-2 (2) (s-I tool-2 (2)

13.6 5.02 4.40 4.61 4.31 -+0.20 NaC10 4 18.6 5.00 3.98 5.26 5.75 + 0.24 0.1 tool (-1 24.9 4.95 4.50 8.32 8.32___0.42 (acetate) 34.9 5.01 4.17 14.4 14.1 +0.70

13.5 4.69 4.33 1.07 1.05 + 0.05 NaCIO 4 17.8 4.70 3.99 1.64 1.66 +__ 0.08 1 tool (-1 25.0 4.69 4.47 2.74 2.80+0.14 (acetate) 35.1 4.69 4.19 6.38 6.28+0.31

13.9 5.03 4.22 4.40 4.34_+0.21 NaCIO 4 19.4 5.01 4.47 5.60 5.53 +0.26 1 tool ( - 1 24.8 4.98 4.36 6.68 6.96 + 0.34 (phosphate) 34.6 5.08 4.48 10.6 10.3 +0.53


Table 5. Temperature effects in NaCI medium. The third-order rate constant (k) is a mean value of at least three experimental values of kow The constant knt is the third-order rate


Medium I Temp. [HSO 3] x 10 a k x 10- 7 knt x 10- ~

(buffer) (°C) - I o g [ H +] (tool d -1) (s -1 tool -2 d 2) (s -1 tool -2 d ~)

13.2 4.99 4.55 7.09 6.13+0.31 NaCI 17.3 5.01 4.04 6.65 7.264-0.35 0.1 mol : - 1 21.3 5.00 4.47 8.60 8.53 4-0.43 (acetate) 24.8 5.00 4.06 8.35 9.80+0.50

34.8 5.00 4.49 15.90 14.34-0.7

13.2 5.00 4.30 7.62 7.424-0.39 NaC1 18.1 5.01 3.80 8.26 8.41 4-0.45 1 mol : - t 24.9 5.00 4.28 9.61 9.904-0.50 (acetate) 34.8 5.01 3.89 12.7 12.44-0.65

Table 6. Temperature effects in NazSO 4 medium. The third-order rate constant (k) is a mean value of at least three experimental values of kob ,. The constant knt is the third-order rate


Medium I Temp. [HSO~] x 10 a k x 10 -7 kfl ' × 10 -~

(buffer) (°C) - I o g [ H +] (mold -1) (s -1 tool -2 d 2) (s -1 mo1-2 d 2)

13.4 5.11 4.01 1.25 1.304-0.06 Na2SO 4 18.6 5.07 4.33 1.68 1.54+0.08 1 mol g - i 24.8 5.02 4.20 1.76 1.864-0.10 (acetate) 34.8 4.87 4.46 2.51 2.49+0.12

Table 7. Activation energy, enthalpy and entropy of activation in different media

Medium I E, AH~ AS:[: (buffer) (mol : -1) (kJ mo1-1) logA (kJ mo1-1) (J K -1 mo1-1)

NaCIO 4 1.0 61. l _+ 4.0 18.1 4- 0.3 58.6 4- 4.0 94 4- 20 (acetate) 0.1 40.3 4- 3.5 15.0 4- 0.4 37.7 4- 3.5 33 4- 6

NaCIO 4 (phosphate) 1.0 30.7 4- 2.5 13.2 + 0.4 28.2 4- 2.5 0 + 6

NaCI 1.0 17.6 4-1.5 11.l -t-0.3 14.9 4-1.5 - 4 3 + 7 (acetate) 0.1 28.7 4:3.0 13.0 4- 0.6 23.9 4- 3.0 - 12 4- 4

Na2SO4 (acetate) 1.0 22.4 4- 2.0 11.24- 0.4 19.9 4- 2.0 -- 39 4- 7

values are given in Table 7. Table 7 shows important variations of the activation energy (factor 3.5).

For the various supporting electrolytes, the values of the activation energy are very similar at low ionic strength (0.1 mol : - 1 ) , but become substantially different when the ionic strength increases (1 mol d-1). Higher values of the activation energy are observed in NaCIO4 medium giving low interactions with reactants. In chloride, sulphate and phosphate media, these species may interact with the reacting species; the activation energy is lower than in perchlorate medium and decreases when the ionic strength increases.

On the other hand, negative or low values of the entropy of activation (AS:~) imply that the intermedi-

ate complex has a lower entropy than the reactants. This case is observed for phosphate, chloride and sulphate; these species seem to interfere in the first pre-equilibria of the mechanism.

Effects of metal ion traces

Many investigations have been carried out on the aqueous-phase oxidation of dissolved SO2 by dissolved oxygen when catalysed by metal ions. However there is a lack of information on the effect of metal ion catalysis on the aqueous-phase oxidation of dissolved SO2 by H202. Results obtained by Mart in (1984) show no such effects. Hoffmann and Jacob (1984) reported important catalytic effects of Fe 2+ and


Table 8. Effects of metal ions on the third-order rate constant

[Metal ion] k(Mn 2+) x 10 -7 k(Cu 2+) × 10 -7 k(Fe a+) x 10 -7 k(Cr 3+) x 10 -7 (mol f - l ) (s- i tool-2 t,2) (s-1 tool-2 f2) (s-I mol-2 ~2) (s-1 mol-2 t,2)

0 8.40 8.40 8.40 8.40 2 × 10 -6 8.43 8.32 10.3 9.44 1 x 1-0 5 8.03 8.22 9.23 8.73 5 x 10 -5 8.16 8.44 8.33 9.18 1 x 10 -4 8.14 9.25 9.26 8.25 2 × 10 -4 8.26 7.62 10.1

Cu 2+. A similar effect for Fe 2÷ was obta ined by Ibusuki et al. (1990).

The kinetics of oxidat ion of dissolved SO2 by H202 were studied in the presence of Cu 2÷, M n 2÷, Fe 3+, Cr 3+ and Fe 2÷ (concentrat ions of 1 0 - 6 - 1 0 -4 mol

f t). NaCIO4 at 0.1 mol f - ~ was the used supporting electrolyte. Firs t -order kinetics were found with respect to hydrogen sulphite, H ÷ ion and hydrogen peroxide. The th i rd-order rate cons tan t (k) was calculated and the values were found to be the same without metal ions and with Cu 2÷, M n 2÷, Fe 3÷ and Cr 3+ ions at various concentra t ions (Table 8). The S(IV) - H 2 0 2 system appears to be insensitive to pos- sible catalytic effects of these metal ions under our experimental conditions. Only Fe 2÷ ion showed a catalytic effect. The var ia t ion of the th i rd-order rate cons tant as a function of the concent ra t ion of Fe 2 ÷ is given in Fig. 6.

The th i rd-order rate cons tan t observed in the presence of the various Fe 2 + concent ra t ions shows a small catalytic effect and may be expressed by a hyperbolic function:

k = ( A + B [ F e 2 + ] ) / ( I + C [ F e 2 + ] ) (5)

k =(8.41 × 107+ 1.25 × 10141-Fe2+])/

(1 +8.25 × 105[Fe+]) .

Fen ton (1894) reported that the ferrous ion strong- ly promotes oxidat ion reactions by hydrogen peroxide. In the literature, al ternative in terpreta t ions of H 2 0 2 - F e 2+ reactions are postulated: intermediate steps with format ion of an OH" radical or a ferryl ion (FeO 2 +) with format ion of ferryl complexes (Rush and Koppenol , 1986; Rahhal and Richter, 1988; Wall- ing, 1974).

Here we propose the hypothesis of a mechanism with a second activated complex between sulphite and FeO 2 + ion. The reaction mechanism occurs in two

steps:

Fast pre-equilibrium:

K

H S O 3 + H 2 0 2 , , in te rmedia te 1

(peroxosulphite)

K

Intermediate 1 + Fe 2 ÷ ~ .' in termediate 2

(FeO 2 + complex).

15-

k

(107M'2~ 1)

11,

+

0.1 1 10 100 1000 10000

Fe(ll) concentration (10"7mol I "1)

Fig. 6. influence of the concentration of Fe 2 + ion on the third-order rate constant. + Experimental data at - l o g [H + ] = 5.0; * experimental data at - l o g [H +] =4.5; the

curve is calculated using Equation (5).

Specific acid catalysis (determining steps):

intermediate 1 + H + ~ sulphate

k'

in termediate 2 + H * --, sulphate + Fe 2 ÷.

This would lead to the following rate law:

k K + k ' K K ' I-Fe 2÷ ] v = 1 + K K ' [ H S O ~ ] [Fe 2÷] [ H + ] [ H S O ~ ] [ ' H 2 0 2 ]

with K ,~ K'. When the concent ra t ion of the intermediate 2 is

low, [Fe 2 +] = [Fetotal], our experimental da ta lead to the following values:

k K = 8 . 4 1 x 107 m o l - 2 f2 s -1

k 'K K' = 1.25 x 1014 mol - 3 t~3 s - l

K K ' = 1.83 x l0 a m o l - 2 f2

k' = 6.83 × l0 s mol - 1 f s - l

CONCLUSIONS

Several uncertainties tha t occur in the l i terature concerning the kinetics and the mechanism of oxida-


tion of dissolved SO2 by H202 have been removed in this work.

This laboratory work has been carried out under conditions approaching those in the atmosphere (concentrated solution drops) with controlled parameters. The experimental results show:

• At constant ionic strength, with a low buffer concentration compared with that of the supporting electrolyte, the effect of the buffer on the kinetics of oxidation of sulphite by H202 can be neglected.

• The rate constant is a function of the nature and the concentration of the supporting electrolyte. This effect is important. The influence of ionic strength can explain the difference between the kinetic rate constant determined by laboratory experiments and the same constant measured by field studies. In clouds, water droplets growth around aerosol deliquescent particles. There are several successive phases of condensat ion-evaporat ion during the droplet 's life in clouds. During this period the ionic strength of these droplets is variable and may reach high values.

• Only Fe z+ species appear to be catalysts for the oxidation of SO2 by H 2 0 2.

• The determination of the temperature dependence allows the extrapolation of these data to the various real atmospheric temperature conditions.

Acknowledgements--This work was supported in part by the Commission of the European Communities under contract STEP 0005C, by the 'CNRS PIREN' and by the 'Minist6re de l'Environnement', Paris. This work is a contribution to EUROTRAC subproject HALIPP.

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Electrolyte effects on aqueous atmospheric oxidation of sulphur dioxide by hydrogen peroxide