2450
P. A.
(7) K. S . Bagdasar'yan, Z. A. Sinitsyna, and V. I. Muromtsev, Dokl. Akad. Nauk SSSR, 153,374 (1963). (8) S. Siege1and K. Eisenthal, J. Chem. fhys., 42, 2494 (1965). (9) K. Shimokoshi, Y. Mori, and I. Tanaka, Bull. Chem. SOC. Jim..409 302 (1967). (IO) J. E. Bennett and 6. Mile, Trans. faraday Soc., 67, 1587 (1971). (1 1) P. J. Krusic and T. A. Rettig, J. Am. Chem. Soc., 92, 722 (1970). I. SOC. Jpn.. 45, 1357 (1972). (12) S. Arlmitsu and H. Tsubomura, &/Chem.
Skotnlchi, A. G.Hopkins, and C. W. Brown
(13) E. Hayon, T. Ibata, N. N. Lichtin, and M. Simic, J. fhys. Chem., 76, 2072 (1972). (14) M. C. R. Symons and M. Townsend, J. Chem. fhys.. 25, 1299 (1956). (15) S. L. Murov, "Handbook of Photochemistry", Marcel Dekker, New York, N.Y., 1973. (16) A. G. Harrison and F. P. Lossing, Can. J. Chem., 37, 1696 (1959). (17) R. Wilson, J. Chem. SOC.E, 84 (1968). (18) N. Filipescu and F. L. Minn, J. Am. Chem. Soc., 90, 1544 (1968).
Time Dependence of Quantum Yields for the Photooxidation of Sulfur Dioxide Peggy A. Skotnlcki, Alfred 0. Hopkins, and Chris W. Brown* Department of Chemlstty, University of Rhode Island, Klngston, Rhode Island 0288 1 (Received August 9. 1974; Revised Manuscript Received June 30, 1975)
Quantum yields for the photooxidation of SO2 have been studied in a static system as a function of time using Raman bands of SO2 and so3 to determine the extent of reaction. The quantum yields were found to decrease during the first 5 to 10 hr. It ie suggested that this decrease is due both to the formation of a film on the walls of the reaction cell and to a back reaction.
Most previous inve~tigationsl-~ of the photooxidation of SO2 in its first allowed absorption region (2600-3300 A) were performed at fixed times, since wet chemical analysis made it difficult to do time-dependent studies. The amount of SO3 formed from SO2 was determined by dissolving the SO3 in a suitable solution, and measuring the amount of S042- formed. Cox1 showed that this method could lead to serious errors if SO2 dissolved forming S032-, which could be converted to S042- in the solution. In order to avoid this problem he ran blanks with pure SO2 to determine the amount of S042- formed from S02. Recently, we developed a new method for the analysis of an so3-so2 gas m i ~ t u r e Both .~ SO2 and SO3 have one strong band in their Raman spectra: 1150 cm-l for SO2 and 1068 cm-I for SO3. The spectrum of a gaseous mixture can be measured without removing the gases from the reaction cells. From the intensities of the two bands and calibration data on intensity vs. pressure, the amounts of SO2 and SO3 present can be determined. It takes less than 10 min to measure the Raman spectrum in the region of interest; thus, it is possible to do time dependent studies of the photooxidation of S02, i.e., to remove the reaction cell from the photolysis chamber, measure the Raman spectrum, and return the vessel to the chamber for continued photolyais. Photolysis experiments were carried out in a Rayonet Srinivasan-Griffen photochemical reactor with lamps emitting a band centered at 3000 A. A 35-cm long, 4-cm diameter cylindrical tube (both quartz and Pyrex were used) with a Teflon stopcock at one end was used as the reaction cell. The cell was cleaned prior to each use by soaking in hot chromic acid overnight, rinsing with 3 N aqueous ammonia overnight, and washing with distilled water and deionized water several times. It was then evacuated to Torr, and degassed by heating to -300'C. Finally, a small pressure of oxygen was introduced and subjected to a microwave discharge to oxidize any impurities remaining on the walls. SO2 (Matheson, 99.98%) and oxygen (M. G. ScientifThe Journal of Physical Chemlstry, Vo/. 79, No. 22, 1975
ic, 99.99%) were vacuum distilled several times prior to use and the reaction cell was filled using standard manometric procedures. Quantum yields for SO3 were determined from the amount of SO3 formed and intensity of the light absorbed ( I , ) by S02. The pressure of so3 relative to SO2 was determined from the Raman ~ p e c t r a this ; ~ was converted to absolute pressures using the known initial pressure of SO2 and assuming that for each molecule of SOB formed two molecules of SO2 reacted. The incident intensity was determined by a diethyl ketone actinometer for both quartz and Pyrex cells. Extinction coefficients for both diethyl ketone and SO2 were estimated from published ~ p e c t r a ,and ~ it was assumed that absorptions of both substances followed Beer's law. Raman spectra were measured on a Spex Industries Model 1401 double monochromator, with a CRL Model 52-A argon ion laser (-600 mW at the sample) and photon counting detection. The spectral measurements were made with the 4880-A laser line but, in several experiments, the results were confirmed with the 5145-A laser line to eliminate the possibility of surface films absorbing in the region of Raman scattering. Details of the measurements and calibration data are discussed el~ewhere.~ Cox1 pointed out that the major problem with all of the previous studies of the photolysis of SO2 has been the wide range of quantum yields obtained. In this note we show evidence for these discrepancies. In Figure l the quantum yield for SO3 is plotted vs. time for an initial SO2 pressure of 200 Torr. After 1 hr, -1.5% of the SO2 is converted giving a quantum yield for SO3 of (7.4 f 0.7) X After 3 hr a conversion of -2.4% SO2 is reached with a quantum yield for SO3 of (4.1 f 0.4) X The quantum yield levels off after about 15 hr at -0.5 X lov3. The quantum yield for SO3 vs. time for an initial SO2 pressure of 100 Torr is plotted in Figure 2. The results are practically the same as for the initial SO2 pressure of 200
2451
Photooxidation of Sulfur Dioxide
'i
mX 2
-
6
0
20 Time, h r s
Figure 1. Quantum yields for SO3 vs. time. Pure SO2 at an initial pressure of 200 Torr in a quartz reaction cell.
8
c
2tL bL
0
IO Time,
20 hrs
Figure 2. Quantum yields for SO3 vs. time in a quartz reaction cell: (A)pure SO2 at an initial pressure of 100 Torr; (0) SO2-02 mixture at an initial pressure of 102 and 52 Torr, respectively.
Torr, i.e., the quantum yield starts at a maximum and falls with time to a minimum level (-0.5 X at -15 hr. In a Pyrex cell similar results are obtained (Figure 3). The decrease of the quantum yield for so3 with time could be the reason for the apparent discrepancies in the published data. Previous experiments were performed both in static and flow systems for different time periods and, since $sos depends upon time, this would explain the range of results. There are two possible explanations for the apparent decrease in the quantum yield of SO3 in a static system; condensation of a film on the vessel walls and back reaction. Other investigators do not report the visible appearance of condensation on the walls; however, the photolysis of SO2 must lead to the formation of a lower oxide of sulfur or pure sulfur, both of which might condense on the vessel walls. We could not see any visible fogging during the first few hours of photolysis, although $sos did decrease during this time. However, fogging of the cell walls was noticed after several hours of photolysis when the cell was placed in the laser beam for the Raman measurements. After long periods of photolysis fogging became visible in room light. To determine if condensation decreased the quantum yields of SO3, we photolyzed SO2 (200 Torr) for 24 hr in a small cylindrical quartz cell, which had optically flat windows. After photolysis a film was visible on the walls. The cell was evacuated to Torr for 5 min, and then the uvvisible spectrum of the film was measured. The spectrum showed that a considerable amount of SO2 remained in the cell even though it had been evacuated. The SO2 could have been absorbed by the film or formed from the film. After evacuating the cell for 15 hr at Torr most of the film visibly disappeared; however, the cell still absorbed below 3300 A suggesting that a film was present. Since the film started to sublime as soon as the photolysis was stopped, it is not possible to give a quantitative value for its effect on the apparent quantum yields. However, our measurements showed that the film reduced the incident intensity by a t least 10%. In a separate experiment we measured the Raman spectrum of the film immediately after photolysis. The spectrum was very similar to rhombic sulfur.6 Bands due to SO groups were not observed; however, it should be noted that sulfur is a very strong Raman scatterer, and the presence of SO groups should not be ruled out. All types of sulfur (rhombic, polymeric, etc.) absorb below 3500 A;7 therefore, sulfur could account for the absorption by the film. Calvert and coworkers2 suggest that the photolysis of SO2 in its first allowed absorption region (2600-3300 A) leads to the following reactions:
SO2 + hu
m
-
0
X
D
8
-
+
SO2*
(1)
+
so2*+ so2 SOB so (2) Very recently, Chung, Calvert, and Bottenheim8 reported on the quantum yield of so3 in static and flow systems. They obtained evidence that the decrease of 4 ~ in0 slow ~ flow and static systems is due to the following back reactions: so + so.? 2 s 0 2 (3)
-
and 2SO
Flgure 3. Quantum yields for SO3 vs. time in a Pyrex reaction cell: (A)pure SO2 at an initial pressure of 100 Torr; (0) SO2-O2 mixture at an initial pressure of 72 and 30 Torr, respectively.
-
SO:!
+ S (or (SO)2)
(4) Both of these reactions are likely in a static system. Presence of the film on the vessel walls and of sulfur in the film The Journal of Physical Chemistry, Vol. 79, No. 22, 1975
Communications to the Editor
2452
definitely supports the latter reaction. Furthermore, evidence of SO2 in the uv spectrum of the “evacuated” cell would also support the existence of a film on the walls; either SO2 was adsorbed by the film or formed from the film by reaction 4. Possibly, SO radicals condense on the walls in the form of a sulfur-oxygen polymer. After the gaseous mixture of sors03 was removed, the polymer decomposed to SO2 and sulfur. This explanation would also be supported by the evidence of sulfur in the Raman spectrum of the film. Condensation on the walls and back reactions should both be reduced by adding oxygen to SO2 prior to photolysis; therefore, we measured the quantum yields for SO2-02 mixtures. For a mixture of 102 Torr of SO2 and 52 Torr of 0 2 in a quartz cell (Figure 2) &03remains constant for -10 hr and then decreases. For a mixture of 72 Torr of SO2 and 30 Torr of 0 2 in a Pyrex cell (Figure 3) $so3 increases slightly during the first few hours and then levels off. The difference between the quartz and Pyrex cells might be explained by the fact that Pyrex cuts off much of the light emitted by the 3000-A lamps. However, the important point is that the quantum yields for SO2-02 mixtures are about the same as those obtained during the first few hours of the photolysis of pure S02. Apparently, oxygen reacts with S02* or SO to give SO3;thus, it inhibits formation of a film on the walls and the possibility of a back reaction between SO and SO3 by reaction 3. The present results confirm the conclusion of Chung,
Calvert, and Bottenheim8 that true quantum yields for SO3 are obtained only in fast flow systems. In the case of a static system, true yields are obtained only a t short photolysis times. In addition, our results strongly indicate that the back reaction in eq 4 (or a very similar reaction) is responsible for reducing the quantum yields. Acknowledgment. We wish to express our appreciation to the National Science Foundation for a matching grant which made possible the purchase of the Raman instrumentation. We also wish to thank Professor Jack G. Calvert for making his results available to us prior to their publication.
References and Notes (1) (2)
R. A. Cox, J. Phys. Chem., 78, 814 (1972). S.Okuda, T. N. Rao, D. H. Siater, and J. G. Calvert, J. Phys. Chem., 73, 4412 (1969); T. N. Rao, S. S. Collier, and J. G. Calvert, J. Am. Chem. Soc., 91, 1609 (1969); T. N. Rao. S. S.Collier, and J. G. Calvert, /bid.,
91, 1616 (1969). (3) T. C. Hall, Jr., Ph.D. Thesis, University of California, Los Angeles, Calif., 1953. (4) P. A. Skotnicki, A. G. Hopkins, and C. W. Brown, Anal. Chem., 45, 2291 (1973). (5) J. G. Calvert and J. N. Pitts, “Photochemistry”. Wiley, New York, N.Y., 1966. (6) R. E. Barletta and C. W. Brown, J. Phys. Chem., 75, 4059 (1971). (7) B. Meyer, M. Gouterman, D. Jensen, T. V. Oommen, K. Spitzer, and T. Stroyer-Hansen in “Sulfur Research Trends”. D. J. Miller and T. K. Wiewiorowski, Ed., American Chemical Society, Washington, D.C.. 1972, pp 53-72. (8) K. Chung. J. 0. Calvert, and J. W. Bottenheim, lnt. J. Chem. Kinet., 7, 161 (1975).
COMMUNICATIONS TO THE EDITOR
Comments on the Paper “Effects of Anions on the Potentials of Zero Charge of Metals” by A. K. Vijh Publication costs assisted by the National Research Council. Rome, ltaly
Sir: In the paper under discussion,’ the author attempts to reconcile existing different view^^-^ on the types of energy involved in specific adsorption of ions at electrodes. The author proposes a formulation derived from the application of a Born-Haber cycle to the reaction of adsorption envisaged as a process of formation of a surface compound. He obtains an equation where the major factors considered in different theories, i.e., surface bond formation energy and ionic hydration energy, respectively, appear together. On one hand, this approach probably improves the qualitative picture of the adsorption process somewhat, especially in view of some experimental evidence for partial charge tran~fer63~ in ionic adsorption. On the other hand, the treatment also shows some conceptual oversights which make the reader unavoidably suspicious toward the whole approach. First of all, in the Born-Haber cycle the author introduces the enthalpy of formation of the bulk compound instead of that of the surface compound on the grounds that the surface bond actually produces a demetallization of the surface. This concept may be found in chemisorption theories8 but it is also true that it can be questioned someThe Journal of Physical Chemistry, Vol. 79, No. 22, 1975
what.g At least,-it cannot be expected to be generally applicable. If demetallization may occur upon extensive oxidation of a metal surface in the gas phase or even electrochemically, such a process can hardly be envisaged in ionic adsorption under electrical control at the potential of zero charge. One of the reasons given by the author to justify the use of the formation heat for the bulk compound is that the Pauling equationlo to evaluate bond strengths applies only to isolated molecules in the gas phase. This is not true in principle in that it has been shown in the case of hydrogen adsorption on transition metals that the initial heat of adsorption can satisfactorily be calculated with the Pauling equation if the surface energy of metals is used in the place of the atomization heat to account for the M-M bond energy, and if the metal surface electronegativity is appropriately expressed through the work function. The limitation in the use of the Pauling equation with sp metals is rather to be sought in the intrinsic band structure of these metals so that it is difficult to envisage dangling orbitals at the surface. Therefore, it should clearly be stated that the use of formation heats for bulk compounds is just a first step in correlations.12 Results apparently show13 that formation heats and adsorption heats are likely to be linearly related, but in fact there is no evidence in the case of sp metals that formation heats could be taken quantitatiuely as adsorption heats. The major criticism regards the physical picture of the
