- Email: [email protected]
The reaction of sulfur dioxide with ozone in water and its possible atmospheric significance
.4tmospherir
Encironmenf
Vol. 11, pp. 813-817
S’ergamon
Press1977. Printed inGreatBritain.
THE REACTION OF SULFUR DIOXIDE WITH OZONE IN WATER AND ITS POSSIBLE ATMOSPHERIC SIGNIFICANCE RONALD
E. ERICKSON, LELAND M. YATES, ROBERT L. CLARK and DAVID MCEWEN
Department of Chemistry, University of Montana, Missoula, Montana 59801, U.S.A. (First received 25 February and in &ml farm 27 December 1976)
Abstract-The rate of the reaction between ozone and dissolved sulfur dioxide is strongly pH dependent. This is so because sulfite ion reacts extremely rapidly (second order rate constant four orders of magnitude higher than that of bisutite ion). These results suggest that under some conditions, atmospheric oxidation of sulfur dioxide may involve ozone.
INTRODUCTION
Sulfur dioxide is known to be converted into sulfate (sulfuric acid) in the atmosphere by a number of chemical processes (Urone and Schroeder, 1969; Bufalini, 1971: Sethi, 1971; Sidebottom, Badcock, Jackson and Calvert, 1972; Wilson, Levy and Wimmer, 1972; Freiberg, 1975). We report here on the chemical kinetics of the reaction of ozone with various sulfur (IV) species in aqueous solutions. A suggestion by R. D. Cadle (1969) that a study of the ozonation of sulfurous acid in fog droplets might be appropriate prompted us to undertake this study. Despite the many kinetic studies carried out since 1969, a recent review by Finieyson and Pitts (1976) still states, “Present un~er~~ties in the rate constants [referring to SO2 oxidation by OH, HOz, 0(3P), Oj, NO,, N,O,, ROz and RO] and in many of the products of these reactions preclude a quantitative assessment of their importance, although OH is probably important.” Sulfur dioxide is readily soluble in water, forming some hydrated species of sulfur dioxide, which in turn ionizes to bisulfite and sulfite ions. SOz(aq)
k% H+ + HSO; K1o,Q, = 1.3 x lo-‘M. (1)
HSO;
+> H’ + SO;2K2(25’, = 1.0 x 10-‘M.
(2) Each of these species is capable, in principle, of reacting with ozone as is shown in equations 3-5, and the relative rate of these possible reactions determines the actual overall reaction path. 03
03
+
+
SO2W)
5
wso;
‘F;
o3 + so;
Hlo 2
;5;
H,SO, HSO; so;
+ OZ. + OZ.
+ 02.
(3) (4) (5)
While this work was ia progress Penkett (1972) reported on the kinetics of reaction 4, that of ozone 813
with bisulfite ion, using a pH of 4.65 and assuming all reaction is due to bisulfite ion, which is the predo~nant species under these conditions. He showed that the reaction is second order, first order in both ozone and bisulfite ion. Our results, extrapolated to his conditions, differ considerably from his in some respects (although the possible significance of the reaction to atmospheric conversion was clearly stated by Penkett at that time), and we will consider those differences after reporting our methodology and results. METHODOLOGY
We determined the rate of ozonation of total sulfur species present in buffered aqueous solutions of sulfur dioxide at variable pH by a stopped flow technique. The stopped flow system used was an Aminco-Morrow stopped ilow spectrophotometer with associated U.V.source and a Biomation transient recorder Model 610 with a Cathode Ray Tube display. The transient recorder storage was then fed into a Bausch & Lomb VOM-10 recorder. This enabled the signal collected over a few milliseconds to be recorded on a 10 second virtual time base for a permanent record. The reaction was followed by observing the decrease in ozone concentration (initial concentrations of ozone = 5 x 10m5 to 5 x 10m4M, total sulfur (IV) = low5 M, pH range of - 1.30 to 4.02, T = 25°C and 16°C). Five traces were recorded with each run and three runs were made with each pair of reactant solutions. The solutions of ozone were prepared by bubbling ozone into the proper buffered solution with nitrogen from ozone saturated silica gel at -78°C and were analyzed spectrophotometrically in the stopped flow apparatus (1 = 276 nm, E = 2150); A = 250 nm, E = 2430). The KI method was used for standardization. The S(IV) solutions were made by dissolving reagent grade sodium sulfite in the desired buffer and adding H,S04 if necessary to give the desired pH.
814
R. E. ERICKSoh. L. M. YATES, R. L. CLARK
Concentrations were determined spectrophotometritally at 270 (6 = 461) 276nm (E = 500). This yielded the concentration of sulfur dioxide directly (two wavelengths gave results within 34:; of one another) and allowed equilibrium calculations for other S(W) species. Some of the more acidic solutions were prepared by passing SO, into the solution of sulfuric acid used for the buffer (cn. 2 M HzS04 for pH -0.30). At extremely high acidities both ozone and sulfur dioxide absorb at 276 nm and proper corrections were made in calculating ozone concentration. At most pH values (pH > 2) the (cont. x E) factor for aqueous sulfur dioxide become less significant. A possible source of experimental error would be the decomposition of ozone caused by water, trace impurities or the buffer. Normal buffer solutions using phthalates or ammonium compounds react rapidly with ozone and were completely unacceptable. However acid phosphate buffer systems did not destroy ozone at an appreciable rate (t1/2 z 3 h) and no ozone loss corrections were necessary for the very fast reactions measured in these experiments. RESULTS AND DISCUSSION
The reaction between ozone and sulfur (IV) species follows second order kinetics. first order in ozone and in sulfur (IV), at every pH which we employed. Rate constants were determined using the following equation obtained by integration of the rate equation (- d[O,]/dt = k[O,] [S(W)] assuming the observed stoichiometry of the reaction : k = [l/a - b][ln(b(a - x)/a(b - x))]/t k = rate constant where a = initial ozone concentration h = initial sulfur concentration .Y = change in concentration of 0, at 1 = time of measurement. Table I shows typical rate data. Similar constancy of data results was found at all pH values. The constancy obtained using the equation for a second order reaction, first order in each reactant. is interpreted to show that the reaction, overall. of ozone with S(IV) obeys this rate law at all pH values studied. Our original supposition was that although all sulfur species might undergo oxidation by ozone in aqueous solution, manipulation of the pH of the reaction could “isolate” reactions of the individual species Table
1. Typical
and D. McEwr~
20
18-
.
16-
14“o
IZ-
Xi ,o_m $ 8-
. /!
./
6-.
4 .-l
2-
.
_/’
.
0
I
PH
Fig. 1. Dependence
of rate on pH for 75 C runs
li is in
(M/l s)- I. (aqueous SO,, bisulfite ion and sulfite ion). For example, using the values of the equilibrium constants shown in Equations 1 and 2 at a pH of -0.30 ([H+] = 2M)
CH+ICHSO,j= SO2W
1,5
x ,.
2
then
and
or the S(IV) species present are in the approximate ratio of 133:1:5 x 10e8 while at a pH of 3([H+] = tom3 M). HSO; SO2 (aq)
so;”
1.5 x 10.’ -~~ 1o-J 1.0 x
HSO, and the ratio 0.07: 1: lo-“.
kinetic
of
lo--
10 3 the
species
= 15
=lxlO~” present
data (pH = 3.12)
t (sl
Total S(W) (b - XI M/l
b(a - x) a(b - x)
In& - .x) a(b - x)
0.0004 0.0008 0.0002 0.0016 o.cQ20
5.89 x 1O-4 4.85 4.08 3.49 3.03 2.65
8.84 x IO-“ 7.80 7.03 6.44 5.98 5.60
1.07 1.15 1.23 1.31 1.41
0.069 0.138 0.207 0.274 0.342
0
I 4
I 3
2
(0 - XI M/l
[%I
/
&w, (M/l s)-
’
5.86 x IO5 5.85 5.84 5.80 5.80
becomes
815
Reaction of sulfur dioxide with ozone in water Figure 1 shows a plot of pH vs k,o,a,s(IM,where ktota,s,,v, represents the total second order rate constant for reactions 3, 4, and 5. The figure shows a stronger dependence of the rate constant on pH than we had anticipated. Even if aqueous SOZ itself has a very low reaction rate, at all pH > 2 and < 6 HSO; is the predominant species present; therefore, if all reactions were with this species, the rate dependence on pH in this range would be small, contrary to our observations. Likewise, if the reaction were due entirely to sulfite ion, the dependence on pH would have been larger than that which was observed. We are left with two alternatives: either [H+] has a strong inhibitive effect on the rate or the concentration of the species reacting is dependent on the reciprocal of the [H+]. We have chosen the latter as the most logical explanation of our observations. Using a series of simultaneous equations including the fractions of the different species present and the observed rate constant at various pH values, we can test our assumption. As noted we observed first order kinetics for both S(IV) and [0,] or second order overall at ail pH values studied. If this observation is valid over a range of pH values, then if our assumption of species reaction rates is correct, we arrive at the conclusion that each species reacts with the same order kinetics, but different rate constants. The following relationship then would hold:
in Table 2. ksc,v~,CWhl = ko,ho,CWVIl~ + km;h&JS(IV)I,
or dividing by [S(IV)], k,,,,
but CSOziasJ= WVT CHSW=
= kso,.fso, + ho,
f bo;&cp
All of the k’s being second order rate constants would have the units of 1 mole-’ set-’ or [MS]-‘. The only real assumption in this treatment is that reactions 1 and 2 are rapid enough under the stated conditions so that equilibrium is reached. Equilibrium constants are temperature dependent and to calculate the fractions of the species present, we used the equations of Johnstone and Leppla for values of K1 at 16°C. No temperature dependence data were found for K2 so we used the same value at 16°C as at 25°C. At the pH of interest, this is justifiable because the effect of a few percent difference in this value would have virtually no effect on the calculated concentration of SO;’ and its fraction of the total S(IV). Using the equations in Table 2 in all combinations, it is possible to calculate specific rate constants as follows:
k,,,,,,,Cs(Wl~O,l = $o,CS%J?h~ + biso;WO;lP,I
+ ho,4o;dXVI~
k,,;(25”C)
=
3.1 T 0.24 x lo5 (M/l s)-’
kso, = (25°C) = 2.2 T 0.35 x 109 (M/l k,so,(l6”C)
sj-’
x.&ox WV), x.&o;
[so;21 = S(Ivh_x&o,2
where f; = fraction of the total S(IV) present as that species. From this we arrive at the relationship shown
(8)
ho, = (16°C) = 1.3 T 0.4 x lo9 (M/l s)-‘.
(9)
for the reaction between ozone and sulfurous acid (or sulfur dioxide dissolved in water) was at least two orders of ma~itude less than that of bisulfite ion and became a si~ificant contribution to the overall rate only at pH < 2. To determine the rate constant exactly would require extensive kinetic runs in highly acidic (negative pH) media. The reaction is obviously of no environmental importance and was not studied.
Table 2. Equations for calculations of specific rate constants 25.0” Runs [Fraction HSO,]
2. 3.
-0.30 0.62 1.20
4. 5. 6. 7. 8. 9. 10.
1.71 2.11 2.50 2.80 3.12 3.55 4.02
1.
2.24 9.82 4.31 1.02 1.63 3.03 4.51 5.97 9.47 1.74
x x x x x x x x x x
lo3 IO3 to4 IO5 IO* lo5 to5 lo5 lo5 lo6
= = = = = = = = = = 16.0”
11. I 2. 13. 14.
0.59 2.13 3.09 3.74
+
[Fraction SO,]
0.0065&J; 0.051 kmc; 0.171 k=c,; 0.402 kHSO; 0.627 km,,; 0.805kmo;
0.891kmo; 0.946k,; 0.979 kmo; 0.994 km,,;
+
ks0;L
ho;
pH
+
3.8 x lo-‘Ok
+ + -I+ + + + + +
2.48 3.14 2.39 9.38 2.95 6.54 1.44 4.03 1.20
x x x x x x x x x
-2 U+‘k:;z IO-‘ksO;’ W6k,;2 10’6k,,~ IO-Sk,,l lo-“k,;~ 10-4k,,* 10-4k,;I 1O-3 kso;z
[Fraction SOJ
4% + + + + + + + + + +
0.993 0.948 0.829 0.60 0.375 0.196 0.110
kso, kw, kso, ks,, ko, &, k,*
0.055&, 0.021 kwl 0.007 ksO,
Runs (K, = 1.6 x lo-*, K, = 1.0 x lo-‘)
8.6 x lo3
=
0.058 k=t,;
1.42 x lo5 2.13 x 10’ 1.43 x i06
= = =
0.683 kmo; 0.951 k,; 0.989 km;
(7)
= 1.7 T 0.17 x lo5 (M/l s)-1
The rate constant
+ k&WX21tW
(6)
+ + + +
3.15 2.76 1.64 7.61
x x x x
1O-8 kso;z 10-5~o;z W4k,;z 10-4!+z
+ + + +
0.94 bcIio, 0.32 ho, 0.048 kw2 0.0’ kso>
R. E. ERICKSON,L. M. YATES, R. L. CLARK and D. McEw~.x
816
Substituting the rate constants of the major species into the Arrhenius equation leads to the following formulas: -
d(HS0;) ---iF--
= 1Ol4 exp - 1 l,OOO/RT(HSO;)(O,) (10)
- ‘F
= 10” exp - 10,.5OO/RT (SO;)(O,
droplets of water, pH = 5.0, and a temperature of 10°C. Using the activation energies obtained in this work, we can extrapolate values for kHSOTand /zso, i of I.1 x IO” (M/l)I’s-’ and 7.4 x IO8 (M/l)-‘s-’ respectively. Using the values of
).
All of our rate constants {Equations 6-9). but particularly those involving sulfite ion, have large variances. Two major factors account for the relatively large uncertainty. First, the equilibrium constants (Equations 1 and 2) are of great importance in the calculations and literature values vary widely. Second, the basic assumption behind using those constants is that ionic equilibria are indeed achieved. However, the rate constant of the ozone-sulfite ion reaction is so large that under some of our conditions (those of reasonably high total S(IV) and high pH), it is questionable whether equilibrium was achieved. Since the rate constants for the second ionization constant of sulfurous acid are apparently unknown, it is impossible to calculate the exact magnitude of the problem. Qualitatively, the result of failure of the maintenance of the equilibrium concentration of SO;’ would be a lower concentration than that assumed. Thus, the true &ojZ would be larger than the calculated value. As a result, we can conclude that our value of k r lo9 is a lower limit for the true value. Thus, at any pH above 3 (i.e. any naturally occurring atmospheric water) mosr of the ozone caused oxidation of sulfur species occurs through sulfite ion even though over 99% of the sulfur (IV) in solution is in the form of bisulfite ion or sulfur dioxide. Q~ii~tively, the im~rtance of the ozone-oxidation of sulfur (IV) in the atmosphere is paradoxical. In areas of very low sulfur dioxide concentration the acidity of any atmospheric water could be low (high pH), but the relative concentration of sulfite ion would be high. On the other hand, highly polluted (SO,) air would have atmospheric water of relatively low pH and high sulfur (IV) concentration. However. equilibrium at such acidities would demand proportionately less sulfite ion. The first approximation is that the rate of ozone caused disappearance of sulfur (IV) is independent of the concentration of sulfur dioxide in the air. For q~ntitative calculations, it is necessary to make some assumptions about the quality and moisture content of a specific air mass in order to assess the importance of the ozone reaction. The most obvious point is that the air must contain liquid water in arder for the ozone reaction to be important as it has been shown that ozone does not react rapidly with sulfur dioxide in the vapor phase. For purposes of comparison, let us assume approximately the same concentrations as Penkett did: SOZ at 0.007 ppm, ozone at 0.05 ppm, 0.1 g per m3 liquid
(H )(HSO;) K, = .-..___I
= 0.0184
602iaql) from arrive SOT ozone
the work of Johnstone and Leppla (1934) we at an HSO; concentration of 4 x 10es M, an concentration of about 4 x lo-’ M and an cont. of lo-’ M. So the rate is: dS(IV) ~-= (1.1 x 105 X 4 x 10-5 + 7.4 X 108 dt x 4 x lo- ?1 x 10 ‘) = 3 x Ill-‘(M:l)s-
‘.
For each m3 we have I x 10. ‘1 of solution so we 3 x 10- ’ ’ moles/mJisec now obtain reacted 01 I x 10~‘mo1es/m3~1~ or 6.4 x 10mhg/m”!h. The starting concentration of 0.007 ppm gives 30pg/m3 so (6.4/X)) = 0.32 or 33”,, per h is reacted. Penkett arrived at a figure of 12.6”,, per h. This diffcrence is partially due to the apparently different values used for K,, and K, (he calculated a concentration of HSOJ = 1 x 10. ‘). Another possible contributing factor may lie in the precision of our activation energy and the consequent uncertainty of the extrapolated values for the specific rate constants. All of the above c~~lculations assume that solution is rapid enough that equilibrium between the various phases and species is maintained. Reaction of ozone with low valent metallic ions in solution would decrease its concentration and il such reactions have very high rate constants, then the significance of S(N) reaction is lessened. We do not have the necessary data to test this alternative hypothesis. Comparing the above rate with rates published fat other oxidation modes. it must he concluded that. for the cited conditions (liquid droplets. presence of ozone at a constant ~0Ilcentrat~onj ozone caused oxidation may be an jmpor~nt contribution to the oxidation of S(N) in the atmosphere. For oxamplc, Brimblecombe and Spedding (1974) using approximately the same total S(N) and H,O(l) values that we assume. calculate the removal of about 3!,, per day by the Fe(N) catalyzed oxidation by atmospheric oxygen, Cheng et uI. (1971) calculate a 2”‘);;ihdecrcasc in sulfur dioxide in natural fog in the presence of manganese salts and Sidebottom clt ul. (1972) show a 1.9x,/h loss of sulfur dioxide photochemically at high humidity.
Reaction of sulfur dioxide with ozone in water Obviously the rates we have determined (as well as the lesser rates found by Penkett (1972)) are large compared to those of other oxidation modes and to actual conversion rates in the atmosphere. Finlayson and Pitts (1976) note that actual conversion takes place at between 5 and 10%/h (Los Angeles) to <0.25%/h (Europe). Again, we must emphasize that ozonation as a mode of conversion may be significant
only under conditions in which droplets of water are present.
Cheng R., Corn M. and Frohliger J. (1971) Contribution to the reaction kinetics of water soluble aerosols and SOZ in air at PPM concentrations. Atmospheric Enuironment 5, 987-1008.
Finlayson B. J. and Pitts J. N. Jr. (1976) Photochemistry of the polluted troposphere. Science 192, 111-119. Freiberg J. (1975) The mechanism of iron catalyzed oxidation of SO* in oxygenated solutions. Atmospheric Environment 9, 661-672.
Johnstone H. F. and Leppla P. W. (1934) The solubility of sulfur dioxide at low partial pressures. The ionization constant and heat of ionization of sulfurous acid. J. Am. them. Sot. 56, 2233-2238.
Acknowledgement-The
Protection research.
817
authors thank the Environmental Agency for a grant which supported this
REFERENCES
Brimblecombe P. and Spedding D. J. (1974) The catalytic oxidation of micromolar aqueous sulfur dioxide. Atmospheric Environment 8. 937-945.
Bulfalini M. (1971) Oxidation of sulfur dioxide in polluted atmospheres--A review. Environ. Sci. Technol. 5, 685-700.
Cadle R. D. (1969) Comments on environmental appraisal: Oxidants, hydrocarbons, and oxides of nitrogen by James N. Pitts. J. Air POW. Control Ass. 19, 668.
Penkett S. A. (1972) Oxidation of SO2 and other atmospheric gases by ozone in aqueous solution. Nature, phys. Sci. 240, 105106. Sethi D. S. (1971) Photo-oxidation of sulfur dioxide. J. Air PoUut. Control Assoc. 21, 418420.
Sidebottom H. W., Badcock C. C., Jackson G. E., Calvert J. C., Remhardt G. W. and Damon E. K. (1972) Photoxidation of sulfur dioxide. Environ. Sci. Technol. 6, 72-79. Urone P. and Schroeder W-(1969) SO, in the Atmosphere: A wealth of monitoring data but few reaction rate studies. Environ. Sci. Technol. 3. 436-145. Wilson W. E. Jr., Levy A., Wimmer D. (1972) A study of sulfur dioxide in photochemical smog II&Effect of sulfur dioxide on oxidant formation in photochemical smog. J. Air Pollut. Control Ass. 22, 27-32.
Encironmenf
Vol. 11, pp. 813-817
S’ergamon
Press1977. Printed inGreatBritain.
THE REACTION OF SULFUR DIOXIDE WITH OZONE IN WATER AND ITS POSSIBLE ATMOSPHERIC SIGNIFICANCE RONALD
E. ERICKSON, LELAND M. YATES, ROBERT L. CLARK and DAVID MCEWEN
Department of Chemistry, University of Montana, Missoula, Montana 59801, U.S.A. (First received 25 February and in &ml farm 27 December 1976)
Abstract-The rate of the reaction between ozone and dissolved sulfur dioxide is strongly pH dependent. This is so because sulfite ion reacts extremely rapidly (second order rate constant four orders of magnitude higher than that of bisutite ion). These results suggest that under some conditions, atmospheric oxidation of sulfur dioxide may involve ozone.
INTRODUCTION
Sulfur dioxide is known to be converted into sulfate (sulfuric acid) in the atmosphere by a number of chemical processes (Urone and Schroeder, 1969; Bufalini, 1971: Sethi, 1971; Sidebottom, Badcock, Jackson and Calvert, 1972; Wilson, Levy and Wimmer, 1972; Freiberg, 1975). We report here on the chemical kinetics of the reaction of ozone with various sulfur (IV) species in aqueous solutions. A suggestion by R. D. Cadle (1969) that a study of the ozonation of sulfurous acid in fog droplets might be appropriate prompted us to undertake this study. Despite the many kinetic studies carried out since 1969, a recent review by Finieyson and Pitts (1976) still states, “Present un~er~~ties in the rate constants [referring to SO2 oxidation by OH, HOz, 0(3P), Oj, NO,, N,O,, ROz and RO] and in many of the products of these reactions preclude a quantitative assessment of their importance, although OH is probably important.” Sulfur dioxide is readily soluble in water, forming some hydrated species of sulfur dioxide, which in turn ionizes to bisulfite and sulfite ions. SOz(aq)
k% H+ + HSO; K1o,Q, = 1.3 x lo-‘M. (1)
HSO;
+> H’ + SO;2K2(25’, = 1.0 x 10-‘M.
(2) Each of these species is capable, in principle, of reacting with ozone as is shown in equations 3-5, and the relative rate of these possible reactions determines the actual overall reaction path. 03
03
+
+
SO2W)
5
wso;
‘F;
o3 + so;
Hlo 2
;5;
H,SO, HSO; so;
+ OZ. + OZ.
+ 02.
(3) (4) (5)
While this work was ia progress Penkett (1972) reported on the kinetics of reaction 4, that of ozone 813
with bisulfite ion, using a pH of 4.65 and assuming all reaction is due to bisulfite ion, which is the predo~nant species under these conditions. He showed that the reaction is second order, first order in both ozone and bisulfite ion. Our results, extrapolated to his conditions, differ considerably from his in some respects (although the possible significance of the reaction to atmospheric conversion was clearly stated by Penkett at that time), and we will consider those differences after reporting our methodology and results. METHODOLOGY
We determined the rate of ozonation of total sulfur species present in buffered aqueous solutions of sulfur dioxide at variable pH by a stopped flow technique. The stopped flow system used was an Aminco-Morrow stopped ilow spectrophotometer with associated U.V.source and a Biomation transient recorder Model 610 with a Cathode Ray Tube display. The transient recorder storage was then fed into a Bausch & Lomb VOM-10 recorder. This enabled the signal collected over a few milliseconds to be recorded on a 10 second virtual time base for a permanent record. The reaction was followed by observing the decrease in ozone concentration (initial concentrations of ozone = 5 x 10m5 to 5 x 10m4M, total sulfur (IV) = low5 M, pH range of - 1.30 to 4.02, T = 25°C and 16°C). Five traces were recorded with each run and three runs were made with each pair of reactant solutions. The solutions of ozone were prepared by bubbling ozone into the proper buffered solution with nitrogen from ozone saturated silica gel at -78°C and were analyzed spectrophotometrically in the stopped flow apparatus (1 = 276 nm, E = 2150); A = 250 nm, E = 2430). The KI method was used for standardization. The S(IV) solutions were made by dissolving reagent grade sodium sulfite in the desired buffer and adding H,S04 if necessary to give the desired pH.
814
R. E. ERICKSoh. L. M. YATES, R. L. CLARK
Concentrations were determined spectrophotometritally at 270 (6 = 461) 276nm (E = 500). This yielded the concentration of sulfur dioxide directly (two wavelengths gave results within 34:; of one another) and allowed equilibrium calculations for other S(W) species. Some of the more acidic solutions were prepared by passing SO, into the solution of sulfuric acid used for the buffer (cn. 2 M HzS04 for pH -0.30). At extremely high acidities both ozone and sulfur dioxide absorb at 276 nm and proper corrections were made in calculating ozone concentration. At most pH values (pH > 2) the (cont. x E) factor for aqueous sulfur dioxide become less significant. A possible source of experimental error would be the decomposition of ozone caused by water, trace impurities or the buffer. Normal buffer solutions using phthalates or ammonium compounds react rapidly with ozone and were completely unacceptable. However acid phosphate buffer systems did not destroy ozone at an appreciable rate (t1/2 z 3 h) and no ozone loss corrections were necessary for the very fast reactions measured in these experiments. RESULTS AND DISCUSSION
The reaction between ozone and sulfur (IV) species follows second order kinetics. first order in ozone and in sulfur (IV), at every pH which we employed. Rate constants were determined using the following equation obtained by integration of the rate equation (- d[O,]/dt = k[O,] [S(W)] assuming the observed stoichiometry of the reaction : k = [l/a - b][ln(b(a - x)/a(b - x))]/t k = rate constant where a = initial ozone concentration h = initial sulfur concentration .Y = change in concentration of 0, at 1 = time of measurement. Table I shows typical rate data. Similar constancy of data results was found at all pH values. The constancy obtained using the equation for a second order reaction, first order in each reactant. is interpreted to show that the reaction, overall. of ozone with S(IV) obeys this rate law at all pH values studied. Our original supposition was that although all sulfur species might undergo oxidation by ozone in aqueous solution, manipulation of the pH of the reaction could “isolate” reactions of the individual species Table
1. Typical
and D. McEwr~
20
18-
.
16-
14“o
IZ-
Xi ,o_m $ 8-
. /!
./
6-.
4 .-l
2-
.
_/’
.
0
I
PH
Fig. 1. Dependence
of rate on pH for 75 C runs
li is in
(M/l s)- I. (aqueous SO,, bisulfite ion and sulfite ion). For example, using the values of the equilibrium constants shown in Equations 1 and 2 at a pH of -0.30 ([H+] = 2M)
CH+ICHSO,j= SO2W
1,5
x ,.
2
then
and
or the S(IV) species present are in the approximate ratio of 133:1:5 x 10e8 while at a pH of 3([H+] = tom3 M). HSO; SO2 (aq)
so;”
1.5 x 10.’ -~~ 1o-J 1.0 x
HSO, and the ratio 0.07: 1: lo-“.
kinetic
of
lo--
10 3 the
species
= 15
=lxlO~” present
data (pH = 3.12)
t (sl
Total S(W) (b - XI M/l
b(a - x) a(b - x)
In& - .x) a(b - x)
0.0004 0.0008 0.0002 0.0016 o.cQ20
5.89 x 1O-4 4.85 4.08 3.49 3.03 2.65
8.84 x IO-“ 7.80 7.03 6.44 5.98 5.60
1.07 1.15 1.23 1.31 1.41
0.069 0.138 0.207 0.274 0.342
0
I 4
I 3
2
(0 - XI M/l
[%I
/
&w, (M/l s)-
’
5.86 x IO5 5.85 5.84 5.80 5.80
becomes
815
Reaction of sulfur dioxide with ozone in water Figure 1 shows a plot of pH vs k,o,a,s(IM,where ktota,s,,v, represents the total second order rate constant for reactions 3, 4, and 5. The figure shows a stronger dependence of the rate constant on pH than we had anticipated. Even if aqueous SOZ itself has a very low reaction rate, at all pH > 2 and < 6 HSO; is the predominant species present; therefore, if all reactions were with this species, the rate dependence on pH in this range would be small, contrary to our observations. Likewise, if the reaction were due entirely to sulfite ion, the dependence on pH would have been larger than that which was observed. We are left with two alternatives: either [H+] has a strong inhibitive effect on the rate or the concentration of the species reacting is dependent on the reciprocal of the [H+]. We have chosen the latter as the most logical explanation of our observations. Using a series of simultaneous equations including the fractions of the different species present and the observed rate constant at various pH values, we can test our assumption. As noted we observed first order kinetics for both S(IV) and [0,] or second order overall at ail pH values studied. If this observation is valid over a range of pH values, then if our assumption of species reaction rates is correct, we arrive at the conclusion that each species reacts with the same order kinetics, but different rate constants. The following relationship then would hold:
in Table 2. ksc,v~,CWhl = ko,ho,CWVIl~ + km;h&JS(IV)I,
or dividing by [S(IV)], k,,,,
but CSOziasJ= WVT CHSW=
= kso,.fso, + ho,
f bo;&cp
All of the k’s being second order rate constants would have the units of 1 mole-’ set-’ or [MS]-‘. The only real assumption in this treatment is that reactions 1 and 2 are rapid enough under the stated conditions so that equilibrium is reached. Equilibrium constants are temperature dependent and to calculate the fractions of the species present, we used the equations of Johnstone and Leppla for values of K1 at 16°C. No temperature dependence data were found for K2 so we used the same value at 16°C as at 25°C. At the pH of interest, this is justifiable because the effect of a few percent difference in this value would have virtually no effect on the calculated concentration of SO;’ and its fraction of the total S(IV). Using the equations in Table 2 in all combinations, it is possible to calculate specific rate constants as follows:
k,,,,,,,Cs(Wl~O,l = $o,CS%J?h~ + biso;WO;lP,I
+ ho,4o;dXVI~
k,,;(25”C)
=
3.1 T 0.24 x lo5 (M/l s)-’
kso, = (25°C) = 2.2 T 0.35 x 109 (M/l k,so,(l6”C)
sj-’
x.&ox WV), x.&o;
[so;21 = S(Ivh_x&o,2
where f; = fraction of the total S(IV) present as that species. From this we arrive at the relationship shown
(8)
ho, = (16°C) = 1.3 T 0.4 x lo9 (M/l s)-‘.
(9)
for the reaction between ozone and sulfurous acid (or sulfur dioxide dissolved in water) was at least two orders of ma~itude less than that of bisulfite ion and became a si~ificant contribution to the overall rate only at pH < 2. To determine the rate constant exactly would require extensive kinetic runs in highly acidic (negative pH) media. The reaction is obviously of no environmental importance and was not studied.
Table 2. Equations for calculations of specific rate constants 25.0” Runs [Fraction HSO,]
2. 3.
-0.30 0.62 1.20
4. 5. 6. 7. 8. 9. 10.
1.71 2.11 2.50 2.80 3.12 3.55 4.02
1.
2.24 9.82 4.31 1.02 1.63 3.03 4.51 5.97 9.47 1.74
x x x x x x x x x x
lo3 IO3 to4 IO5 IO* lo5 to5 lo5 lo5 lo6
= = = = = = = = = = 16.0”
11. I 2. 13. 14.
0.59 2.13 3.09 3.74
+
[Fraction SO,]
0.0065&J; 0.051 kmc; 0.171 k=c,; 0.402 kHSO; 0.627 km,,; 0.805kmo;
0.891kmo; 0.946k,; 0.979 kmo; 0.994 km,,;
+
ks0;L
ho;
pH
+
3.8 x lo-‘Ok
+ + -I+ + + + + +
2.48 3.14 2.39 9.38 2.95 6.54 1.44 4.03 1.20
x x x x x x x x x
-2 U+‘k:;z IO-‘ksO;’ W6k,;2 10’6k,,~ IO-Sk,,l lo-“k,;~ 10-4k,,* 10-4k,;I 1O-3 kso;z
[Fraction SOJ
4% + + + + + + + + + +
0.993 0.948 0.829 0.60 0.375 0.196 0.110
kso, kw, kso, ks,, ko, &, k,*
0.055&, 0.021 kwl 0.007 ksO,
Runs (K, = 1.6 x lo-*, K, = 1.0 x lo-‘)
8.6 x lo3
=
0.058 k=t,;
1.42 x lo5 2.13 x 10’ 1.43 x i06
= = =
0.683 kmo; 0.951 k,; 0.989 km;
(7)
= 1.7 T 0.17 x lo5 (M/l s)-1
The rate constant
+ k&WX21tW
(6)
+ + + +
3.15 2.76 1.64 7.61
x x x x
1O-8 kso;z 10-5~o;z W4k,;z 10-4!+z
+ + + +
0.94 bcIio, 0.32 ho, 0.048 kw2 0.0’ kso>
R. E. ERICKSON,L. M. YATES, R. L. CLARK and D. McEw~.x
816
Substituting the rate constants of the major species into the Arrhenius equation leads to the following formulas: -
d(HS0;) ---iF--
= 1Ol4 exp - 1 l,OOO/RT(HSO;)(O,) (10)
- ‘F
= 10” exp - 10,.5OO/RT (SO;)(O,
droplets of water, pH = 5.0, and a temperature of 10°C. Using the activation energies obtained in this work, we can extrapolate values for kHSOTand /zso, i of I.1 x IO” (M/l)I’s-’ and 7.4 x IO8 (M/l)-‘s-’ respectively. Using the values of
).
All of our rate constants {Equations 6-9). but particularly those involving sulfite ion, have large variances. Two major factors account for the relatively large uncertainty. First, the equilibrium constants (Equations 1 and 2) are of great importance in the calculations and literature values vary widely. Second, the basic assumption behind using those constants is that ionic equilibria are indeed achieved. However, the rate constant of the ozone-sulfite ion reaction is so large that under some of our conditions (those of reasonably high total S(IV) and high pH), it is questionable whether equilibrium was achieved. Since the rate constants for the second ionization constant of sulfurous acid are apparently unknown, it is impossible to calculate the exact magnitude of the problem. Qualitatively, the result of failure of the maintenance of the equilibrium concentration of SO;’ would be a lower concentration than that assumed. Thus, the true &ojZ would be larger than the calculated value. As a result, we can conclude that our value of k r lo9 is a lower limit for the true value. Thus, at any pH above 3 (i.e. any naturally occurring atmospheric water) mosr of the ozone caused oxidation of sulfur species occurs through sulfite ion even though over 99% of the sulfur (IV) in solution is in the form of bisulfite ion or sulfur dioxide. Q~ii~tively, the im~rtance of the ozone-oxidation of sulfur (IV) in the atmosphere is paradoxical. In areas of very low sulfur dioxide concentration the acidity of any atmospheric water could be low (high pH), but the relative concentration of sulfite ion would be high. On the other hand, highly polluted (SO,) air would have atmospheric water of relatively low pH and high sulfur (IV) concentration. However. equilibrium at such acidities would demand proportionately less sulfite ion. The first approximation is that the rate of ozone caused disappearance of sulfur (IV) is independent of the concentration of sulfur dioxide in the air. For q~ntitative calculations, it is necessary to make some assumptions about the quality and moisture content of a specific air mass in order to assess the importance of the ozone reaction. The most obvious point is that the air must contain liquid water in arder for the ozone reaction to be important as it has been shown that ozone does not react rapidly with sulfur dioxide in the vapor phase. For purposes of comparison, let us assume approximately the same concentrations as Penkett did: SOZ at 0.007 ppm, ozone at 0.05 ppm, 0.1 g per m3 liquid
(H )(HSO;) K, = .-..___I
= 0.0184
602iaql) from arrive SOT ozone
the work of Johnstone and Leppla (1934) we at an HSO; concentration of 4 x 10es M, an concentration of about 4 x lo-’ M and an cont. of lo-’ M. So the rate is: dS(IV) ~-= (1.1 x 105 X 4 x 10-5 + 7.4 X 108 dt x 4 x lo- ?1 x 10 ‘) = 3 x Ill-‘(M:l)s-
‘.
For each m3 we have I x 10. ‘1 of solution so we 3 x 10- ’ ’ moles/mJisec now obtain reacted 01 I x 10~‘mo1es/m3~1~ or 6.4 x 10mhg/m”!h. The starting concentration of 0.007 ppm gives 30pg/m3 so (6.4/X)) = 0.32 or 33”,, per h is reacted. Penkett arrived at a figure of 12.6”,, per h. This diffcrence is partially due to the apparently different values used for K,, and K, (he calculated a concentration of HSOJ = 1 x 10. ‘). Another possible contributing factor may lie in the precision of our activation energy and the consequent uncertainty of the extrapolated values for the specific rate constants. All of the above c~~lculations assume that solution is rapid enough that equilibrium between the various phases and species is maintained. Reaction of ozone with low valent metallic ions in solution would decrease its concentration and il such reactions have very high rate constants, then the significance of S(N) reaction is lessened. We do not have the necessary data to test this alternative hypothesis. Comparing the above rate with rates published fat other oxidation modes. it must he concluded that. for the cited conditions (liquid droplets. presence of ozone at a constant ~0Ilcentrat~onj ozone caused oxidation may be an jmpor~nt contribution to the oxidation of S(N) in the atmosphere. For oxamplc, Brimblecombe and Spedding (1974) using approximately the same total S(N) and H,O(l) values that we assume. calculate the removal of about 3!,, per day by the Fe(N) catalyzed oxidation by atmospheric oxygen, Cheng et uI. (1971) calculate a 2”‘);;ihdecrcasc in sulfur dioxide in natural fog in the presence of manganese salts and Sidebottom clt ul. (1972) show a 1.9x,/h loss of sulfur dioxide photochemically at high humidity.
Reaction of sulfur dioxide with ozone in water Obviously the rates we have determined (as well as the lesser rates found by Penkett (1972)) are large compared to those of other oxidation modes and to actual conversion rates in the atmosphere. Finlayson and Pitts (1976) note that actual conversion takes place at between 5 and 10%/h (Los Angeles) to <0.25%/h (Europe). Again, we must emphasize that ozonation as a mode of conversion may be significant
only under conditions in which droplets of water are present.
Cheng R., Corn M. and Frohliger J. (1971) Contribution to the reaction kinetics of water soluble aerosols and SOZ in air at PPM concentrations. Atmospheric Enuironment 5, 987-1008.
Finlayson B. J. and Pitts J. N. Jr. (1976) Photochemistry of the polluted troposphere. Science 192, 111-119. Freiberg J. (1975) The mechanism of iron catalyzed oxidation of SO* in oxygenated solutions. Atmospheric Environment 9, 661-672.
Johnstone H. F. and Leppla P. W. (1934) The solubility of sulfur dioxide at low partial pressures. The ionization constant and heat of ionization of sulfurous acid. J. Am. them. Sot. 56, 2233-2238.
Acknowledgement-The
Protection research.
817
authors thank the Environmental Agency for a grant which supported this
REFERENCES
Brimblecombe P. and Spedding D. J. (1974) The catalytic oxidation of micromolar aqueous sulfur dioxide. Atmospheric Environment 8. 937-945.
Bulfalini M. (1971) Oxidation of sulfur dioxide in polluted atmospheres--A review. Environ. Sci. Technol. 5, 685-700.
Cadle R. D. (1969) Comments on environmental appraisal: Oxidants, hydrocarbons, and oxides of nitrogen by James N. Pitts. J. Air POW. Control Ass. 19, 668.
Penkett S. A. (1972) Oxidation of SO2 and other atmospheric gases by ozone in aqueous solution. Nature, phys. Sci. 240, 105106. Sethi D. S. (1971) Photo-oxidation of sulfur dioxide. J. Air PoUut. Control Assoc. 21, 418420.
Sidebottom H. W., Badcock C. C., Jackson G. E., Calvert J. C., Remhardt G. W. and Damon E. K. (1972) Photoxidation of sulfur dioxide. Environ. Sci. Technol. 6, 72-79. Urone P. and Schroeder W-(1969) SO, in the Atmosphere: A wealth of monitoring data but few reaction rate studies. Environ. Sci. Technol. 3. 436-145. Wilson W. E. Jr., Levy A., Wimmer D. (1972) A study of sulfur dioxide in photochemical smog II&Effect of sulfur dioxide on oxidant formation in photochemical smog. J. Air Pollut. Control Ass. 22, 27-32.