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Rate constants at 298 k for the reactions SO+SO+M→(SO)2+M AND SO+(SO)2→SO2+S2O
Volume 76. number 2
CHEhllCAL
RATE CONSTANTS SO+SO+IU-,(SO)2+Xl
PHYSICS
LE-ITERS
l-1 August
1980
AT 298# FOR THE REACTIONS AND SO+(SO)2-,S02+Sz0
John T’. HERRON and Robert E. HLJIE rVnrionn1Bur,-au of Sramfnrds. Ci~enrrcnl Kitretrcs Dir isron. ~~‘aslwrgrorr,DC 32?4, Received
1 December
US.4
1960
The decay of SO generated in the reaction 0 r0CS-r CO+ SO in the absence of O1 has been studied at 298 K from 2 to 8 Torr, using a tubular Row reactor coupled to a mass spectrometer. The products of the decay of SO are S. SO:. S20. and (SO),. The kinetics of tile decay of SO could be accounted for on the basis of the mechanism: SO-SO + hl+ ISOl~+hl(3~, SO+SO+S0~+S 141. so+(so12+ SO>+S:O 151. and SO+products 16). where the latter represents either a wall reaction or a slow bimolecuiar reaction with OCS. On the basis of computer modeling we derire k3 = 1.6x 10”cmLmolV2
5 -’ and k5 = Z x lO’“cm
mol-’
s-‘.
1. Introduction Sulfur monoxide [SO) has been identified as a product of the combustion of organic sulfides, and is therefore expected to be an intermediate in the oxidation of sulfur-containing fossil fuels [I]. It may also be involved in the complex chemistry of the Venusian atmosphere [2]. In this paper we present the results of a study of the production and decay of sulfur
monoxide in a flow reactor. leading to kinetic data on the self-reactions of SO.
2. Experimental The experiments were carried out using a tubular flow reactor coupled to a mass spectrometer as described previously [3]. Sulfur monoxide was produced by reacting atomic oxygen (produced in the absence of 02 by means of the reaction N + NO + NZ + 0) with OCS in a heated pre-reactor at 500 K:
o+ocs+co+so. The hot gases from the pre-reactor
(1)
flowed through a section of tubing wrapped with cooling coils (to bring the temperature back to 298 K) into the main reac322
tor, which was made of 2.5 cm i.d. borosilicate glass. The reactor was about 70cm long and was fitted with a series of reactant inlets. The end of the reactor was defined by the sampling orifice of the mass spectrometer. The decay of sulfur monoxide along the reactor was measured by NO-, titration [A]: SO+N02+SOz+NO
_
(2)
At the same time, the various ions of interest in the mass spectrum were monitored. At each inlet along the reactor, the addition of NO2 not only gives
a quantitative measure of the SO partial pressure at that point, but also quenches the reaction. It is thus possible to relate changes in intensity of peaks in the mass spectrum to the partial pressure of SO at each point of addition of NO2 in the reactor. It was found in initial studies that a film, presumably of sulfur, was deposited on the walls of the reactor. To eliminate this problem, we added a small amount of I-GHs (~2 x 10 -’ ’ mol cmm3) between the pre-reactor and main reactor to scavenge atomic sulfur (k(S+ l-C4Hs) = 2.5 x lOI cm3 mol-’ s-’ [5]). Under these conditions sulfur did not deposit on the walls. Thus all data reported here were taken under conditions of added l-&Ha, and in the absence of 02.
Volume
76. number
2
CHEhlICAL
PHYSICS
I December
LE-l-rERS
1980
3. Results and discussion The products found on reacting atomic oxygen with carbonyl sulfide as detected by mass spectrometry were SO, SOZ, $0, and (SO),. [CO, a primary product of the reaction, was not distinguishable from N2 which was the carrier gas.) Addition of 1-butene
leads to the formation of a product at m/e 88, presumably CJH~S, which indicates that atomic sulfur is also a reaction product. Although SO, SZO, and SO1 are readily identified by mass spectrometry it is not possible to relate the ion intensities to partial pressures. Thus the peak at m/e 48 corresponding to SO can arise from the ionization of neutral SO, or from the dissociative ionization of SO,, SZO or (SO),. Similarly, SOT and $0’ can be formed from SOZ or $0 as well as from the dissociative ionization of (SO)?. On the other hand the peak corresponding to [SO); should
not be complicated by contributions due to dissociative ionization of larger species and should thus be directly proportional to the partial pressure of (SO),. An absolute calibration, however, was not possible. Although there have been various reports of the detection of (SO)Z [6,7], the first positive identification was that of Lovas et al. [8] based on the microwave spectrum. Although the thermodynamic properties of (SO)? are not known, bond strengths for S-S single bonds are in the 30-70 kcal mol-’ range [9], so that the (SO)? species is probably thermally stable under the conditions of our experiments_ From a study of the amount of (SO), formed at a fixed reaction time as a function of the initial SO concentration, we found thar the order of the reaction was about 1.5, indicative of a complex reaction mechanism. Furthermore, the rate of decay of SO was clearly pressure dependent, as can be seen from fig. 1. These observations can be explained on the basis of the following mechanism: SO+SO+M-,(SO)-,+M,
(3)
so+so+so,+s.
(4)
so + (SO)? + soz + szo Reaction (5) was proposed earlier by Schenk and Steudel 161. and explains the observed non-integral order of the reaction.
(5)
-
5
I5
rmc IO. Fig.
I. Decay
of SO (open
20
47i-37
15
I
circles1 and formatlon
of (SO)2
(closed circles) as a function of time in a tubular flow reactor. The dashed and solid lines are computer modeling fits to the experimental data as discussed in the text.
In order to derive rate constants for these reactions, we carried out calculations using the computer
modeling program of Brown [IO]. Initially, values of k3, k4, and k5 were chosen by trial and error to obtain the best fit to the measured SO decay curves. In these calculations we took the initial value of the concentration of (SO)? to be zero. Initial value here refers to the concentration of a species at the first
inlet jet of the reactor, which defines the nominal time “zero” of the system_ In fact the initial value of [(SO)J,=o was not zero, but rather about SO?& of the Snal value of [(SO),]_ To account for the concentration of (SO)? at time zero, we first calculated the mass spectrometric sensitivity of (SO)Z from the computed (SO)? concentration and the measured (SO)Z signal at a single point. We arbitrarily chose to make the calculation for a point corresponding to a total pressure of 8 Torr and a reaction time of 16 ms. From the calculated sensitivity we then calculated the initial (SO), concentration. The modeling calculations were
then repeated several times until a constant value was obtained for [(SO)2],=o and for the sensi!ivity of (SO)z. This value for the sensitivity was then used to fit the data not only at 8 Torr total pressure, but also at total pressures of 2 and 4 TOIT. We found that using these derived values for [(SO),],=, required 323
Volume 76. number 2
CHEMICAL
PHYSICS
small changes (about 20%) in the values used for ks and kJ in the modeling calculation. The results of the modeling calculations for SO using the rate constants k3 = 1.6 X 10” cm’ mol-2s-1, ka = 5 x 10’ cm3 mol-’ s-IV and k5 = 2 x 10’” an3 mol-’ s-l are shown as the dashed lines in fig. I. Obviously at low pressure there is an additional loss process for SO. Addition of a simple first-order loss process SO + products T
(61
with a rate constant kg= 4 s-l (obtained by trial and error) leads to the solid Iine fits of fig. 1 for SO and (SO)t. Reaction (6) could be a simple wall reaction or could correspond to a slow reaction of SO with OCS. Because of the great excess of OCS in the system a bimolecular reaction with a rate constant of about 1 I? cm3 mol-’ s-l would be adequate to account for the additional SO loss at low pressure. On the basis of the computer model fits to both the SO and (SO)z data, we believe that the derived rate constants kJ and ks are accurate to about rtr50%. Further experiments, at higher pressures, wouId reduce the importance of reactions (4) and (6), and lead to more reliable values for k3 and ks. The rate constant derived for reaction (4) is not too reliable. There is however little doubt, based on the limited number of experiments we have made on titrating the atomic sulfur produced in reaction (4) with 1-butene, that the value derived for kg is qualitatively correct. Chung et al. [llj, from a study of the photolysis of SOL, derived a value for the rate constant of the reaction SO + SO + products of about 5 x lOa cm3 mol -I s-’ at a total pressure of 25 Ton of SO2. Our data would predict an equivaIent bimolecular rate constant at this pressure of 2.5 x 10” cm3 mol-’ s-l. If the reaction were indeed as slow as reported by Chung et al. [ll], we would not have observed any significant decay of SO in our experiments. 4.
Concfusions
The recombination reaction of sulfur monoxide leads to the formation of a relatively long-liVe,ed (SO)*
324
LETTERS
1 December 1980
intermediate, which is lost primarily through subsequent reactions with 50 leading to SO2 and SZO. The formation of [St& from SO, which has not previously been treated as an important process in SO chemistry, may lead to significant complications in the interpretation of the chemical kinetics of SO. Stedman et al. [K?] have recently reported rate data on reactions of atoms and small radicals with SzO. The complete interpretation of the chemical kinetics of SO (i.e. reactions of SO with 0, 01, etc.) will require similar data on atomic and free radical reactions involving (SO)z.
Acknowtedgement This work was supported in part by the Planetary Atmospheres Program of the National Aeronautics and Space Administration.
References A. Levy, EL. Merryman and W.T. Reid, Environ. Sci. Technol. 4 (1970) 653. R.G. Prinn. Geaphys. Res. Letters 6 (1979) 807. R.E. Huie and J.T. Herran, Intern. J. Chem. Kinetics 5 (19731 197. M.A.A. Clyne. C.J. Hslstead and B.A. Thrush. Proc. Roy. Sot. A295 (1966) 355. R+B. Klemm and D.D. Davis. Intern. J, Chem. Kinetics 5 (1973) 375. P.W. Schenk and R. Steudel, Angew. Chem. Intern. Ed. 4 (1965) 402. R.G.W. Norrish and G.A. Oldershaw, Proc. Roy. Sot. A249 (1958) 249. F,P. Lovas, E, Ticmann and D.R. Johnson. J. Chem. Phys. 60 (1974) 5005. 5.W. Benson, Chem. Rev. 78 (1978) 23. R.L. Brown, A Computer Program for Solving Systems of Chemical Rate Equations, National Bureau of Standards, NBSIR 76-1055 (US Govt. Printing Office, Washington, 1978). K. Chung, J.G. Calvcrt and J.W. Bottenheim. Intern. J. Chem. Kinetics 7 (1975) 161. D.H. Sfedman, H. AIvord and A. Baker-Blocker. J. Phys. Chem. 78 (1974) 1248.
CHEhllCAL
RATE CONSTANTS SO+SO+IU-,(SO)2+Xl
PHYSICS
LE-ITERS
l-1 August
1980
AT 298# FOR THE REACTIONS AND SO+(SO)2-,S02+Sz0
John T’. HERRON and Robert E. HLJIE rVnrionn1Bur,-au of Sramfnrds. Ci~enrrcnl Kitretrcs Dir isron. ~~‘aslwrgrorr,DC 32?4, Received
1 December
US.4
1960
The decay of SO generated in the reaction 0 r0CS-r CO+ SO in the absence of O1 has been studied at 298 K from 2 to 8 Torr, using a tubular Row reactor coupled to a mass spectrometer. The products of the decay of SO are S. SO:. S20. and (SO),. The kinetics of tile decay of SO could be accounted for on the basis of the mechanism: SO-SO + hl+ ISOl~+hl(3~, SO+SO+S0~+S 141. so+(so12+ SO>+S:O 151. and SO+products 16). where the latter represents either a wall reaction or a slow bimolecuiar reaction with OCS. On the basis of computer modeling we derire k3 = 1.6x 10”cmLmolV2
5 -’ and k5 = Z x lO’“cm
mol-’
s-‘.
1. Introduction Sulfur monoxide [SO) has been identified as a product of the combustion of organic sulfides, and is therefore expected to be an intermediate in the oxidation of sulfur-containing fossil fuels [I]. It may also be involved in the complex chemistry of the Venusian atmosphere [2]. In this paper we present the results of a study of the production and decay of sulfur
monoxide in a flow reactor. leading to kinetic data on the self-reactions of SO.
2. Experimental The experiments were carried out using a tubular flow reactor coupled to a mass spectrometer as described previously [3]. Sulfur monoxide was produced by reacting atomic oxygen (produced in the absence of 02 by means of the reaction N + NO + NZ + 0) with OCS in a heated pre-reactor at 500 K:
o+ocs+co+so. The hot gases from the pre-reactor
(1)
flowed through a section of tubing wrapped with cooling coils (to bring the temperature back to 298 K) into the main reac322
tor, which was made of 2.5 cm i.d. borosilicate glass. The reactor was about 70cm long and was fitted with a series of reactant inlets. The end of the reactor was defined by the sampling orifice of the mass spectrometer. The decay of sulfur monoxide along the reactor was measured by NO-, titration [A]: SO+N02+SOz+NO
_
(2)
At the same time, the various ions of interest in the mass spectrum were monitored. At each inlet along the reactor, the addition of NO2 not only gives
a quantitative measure of the SO partial pressure at that point, but also quenches the reaction. It is thus possible to relate changes in intensity of peaks in the mass spectrum to the partial pressure of SO at each point of addition of NO2 in the reactor. It was found in initial studies that a film, presumably of sulfur, was deposited on the walls of the reactor. To eliminate this problem, we added a small amount of I-GHs (~2 x 10 -’ ’ mol cmm3) between the pre-reactor and main reactor to scavenge atomic sulfur (k(S+ l-C4Hs) = 2.5 x lOI cm3 mol-’ s-’ [5]). Under these conditions sulfur did not deposit on the walls. Thus all data reported here were taken under conditions of added l-&Ha, and in the absence of 02.
Volume
76. number
2
CHEhlICAL
PHYSICS
I December
LE-l-rERS
1980
3. Results and discussion The products found on reacting atomic oxygen with carbonyl sulfide as detected by mass spectrometry were SO, SOZ, $0, and (SO),. [CO, a primary product of the reaction, was not distinguishable from N2 which was the carrier gas.) Addition of 1-butene
leads to the formation of a product at m/e 88, presumably CJH~S, which indicates that atomic sulfur is also a reaction product. Although SO, SZO, and SO1 are readily identified by mass spectrometry it is not possible to relate the ion intensities to partial pressures. Thus the peak at m/e 48 corresponding to SO can arise from the ionization of neutral SO, or from the dissociative ionization of SO,, SZO or (SO),. Similarly, SOT and $0’ can be formed from SOZ or $0 as well as from the dissociative ionization of (SO)?. On the other hand the peak corresponding to [SO); should
not be complicated by contributions due to dissociative ionization of larger species and should thus be directly proportional to the partial pressure of (SO),. An absolute calibration, however, was not possible. Although there have been various reports of the detection of (SO)Z [6,7], the first positive identification was that of Lovas et al. [8] based on the microwave spectrum. Although the thermodynamic properties of (SO)? are not known, bond strengths for S-S single bonds are in the 30-70 kcal mol-’ range [9], so that the (SO)? species is probably thermally stable under the conditions of our experiments_ From a study of the amount of (SO), formed at a fixed reaction time as a function of the initial SO concentration, we found thar the order of the reaction was about 1.5, indicative of a complex reaction mechanism. Furthermore, the rate of decay of SO was clearly pressure dependent, as can be seen from fig. 1. These observations can be explained on the basis of the following mechanism: SO+SO+M-,(SO)-,+M,
(3)
so+so+so,+s.
(4)
so + (SO)? + soz + szo Reaction (5) was proposed earlier by Schenk and Steudel 161. and explains the observed non-integral order of the reaction.
(5)
-
5
I5
rmc IO. Fig.
I. Decay
of SO (open
20
47i-37
15
I
circles1 and formatlon
of (SO)2
(closed circles) as a function of time in a tubular flow reactor. The dashed and solid lines are computer modeling fits to the experimental data as discussed in the text.
In order to derive rate constants for these reactions, we carried out calculations using the computer
modeling program of Brown [IO]. Initially, values of k3, k4, and k5 were chosen by trial and error to obtain the best fit to the measured SO decay curves. In these calculations we took the initial value of the concentration of (SO)? to be zero. Initial value here refers to the concentration of a species at the first
inlet jet of the reactor, which defines the nominal time “zero” of the system_ In fact the initial value of [(SO)J,=o was not zero, but rather about SO?& of the Snal value of [(SO),]_ To account for the concentration of (SO)? at time zero, we first calculated the mass spectrometric sensitivity of (SO)Z from the computed (SO)? concentration and the measured (SO)Z signal at a single point. We arbitrarily chose to make the calculation for a point corresponding to a total pressure of 8 Torr and a reaction time of 16 ms. From the calculated sensitivity we then calculated the initial (SO), concentration. The modeling calculations were
then repeated several times until a constant value was obtained for [(SO)2],=o and for the sensi!ivity of (SO)z. This value for the sensitivity was then used to fit the data not only at 8 Torr total pressure, but also at total pressures of 2 and 4 TOIT. We found that using these derived values for [(SO),],=, required 323
Volume 76. number 2
CHEMICAL
PHYSICS
small changes (about 20%) in the values used for ks and kJ in the modeling calculation. The results of the modeling calculations for SO using the rate constants k3 = 1.6 X 10” cm’ mol-2s-1, ka = 5 x 10’ cm3 mol-’ s-IV and k5 = 2 x 10’” an3 mol-’ s-l are shown as the dashed lines in fig. I. Obviously at low pressure there is an additional loss process for SO. Addition of a simple first-order loss process SO + products T
(61
with a rate constant kg= 4 s-l (obtained by trial and error) leads to the solid Iine fits of fig. 1 for SO and (SO)t. Reaction (6) could be a simple wall reaction or could correspond to a slow reaction of SO with OCS. Because of the great excess of OCS in the system a bimolecular reaction with a rate constant of about 1 I? cm3 mol-’ s-l would be adequate to account for the additional SO loss at low pressure. On the basis of the computer model fits to both the SO and (SO)z data, we believe that the derived rate constants kJ and ks are accurate to about rtr50%. Further experiments, at higher pressures, wouId reduce the importance of reactions (4) and (6), and lead to more reliable values for k3 and ks. The rate constant derived for reaction (4) is not too reliable. There is however little doubt, based on the limited number of experiments we have made on titrating the atomic sulfur produced in reaction (4) with 1-butene, that the value derived for kg is qualitatively correct. Chung et al. [llj, from a study of the photolysis of SOL, derived a value for the rate constant of the reaction SO + SO + products of about 5 x lOa cm3 mol -I s-’ at a total pressure of 25 Ton of SO2. Our data would predict an equivaIent bimolecular rate constant at this pressure of 2.5 x 10” cm3 mol-’ s-l. If the reaction were indeed as slow as reported by Chung et al. [ll], we would not have observed any significant decay of SO in our experiments. 4.
Concfusions
The recombination reaction of sulfur monoxide leads to the formation of a relatively long-liVe,ed (SO)*
324
LETTERS
1 December 1980
intermediate, which is lost primarily through subsequent reactions with 50 leading to SO2 and SZO. The formation of [St& from SO, which has not previously been treated as an important process in SO chemistry, may lead to significant complications in the interpretation of the chemical kinetics of SO. Stedman et al. [K?] have recently reported rate data on reactions of atoms and small radicals with SzO. The complete interpretation of the chemical kinetics of SO (i.e. reactions of SO with 0, 01, etc.) will require similar data on atomic and free radical reactions involving (SO)z.
Acknowtedgement This work was supported in part by the Planetary Atmospheres Program of the National Aeronautics and Space Administration.
References A. Levy, EL. Merryman and W.T. Reid, Environ. Sci. Technol. 4 (1970) 653. R.G. Prinn. Geaphys. Res. Letters 6 (1979) 807. R.E. Huie and J.T. Herran, Intern. J. Chem. Kinetics 5 (19731 197. M.A.A. Clyne. C.J. Hslstead and B.A. Thrush. Proc. Roy. Sot. A295 (1966) 355. R+B. Klemm and D.D. Davis. Intern. J, Chem. Kinetics 5 (1973) 375. P.W. Schenk and R. Steudel, Angew. Chem. Intern. Ed. 4 (1965) 402. R.G.W. Norrish and G.A. Oldershaw, Proc. Roy. Sot. A249 (1958) 249. F,P. Lovas, E, Ticmann and D.R. Johnson. J. Chem. Phys. 60 (1974) 5005. 5.W. Benson, Chem. Rev. 78 (1978) 23. R.L. Brown, A Computer Program for Solving Systems of Chemical Rate Equations, National Bureau of Standards, NBSIR 76-1055 (US Govt. Printing Office, Washington, 1978). K. Chung, J.G. Calvcrt and J.W. Bottenheim. Intern. J. Chem. Kinetics 7 (1975) 161. D.H. Sfedman, H. AIvord and A. Baker-Blocker. J. Phys. Chem. 78 (1974) 1248.