4412
NOTES
Table I: Values of (1 - R)/l/-,NacI+mNaC1 as a Function of ~ N ~ for O Iboth the Isopiestic and the Emf Measurements of Sodium Chloride and Calcium Chloride Solutions
0.080 0.160 0.240 0.320
0.0064 0.0256 0.0576 0.1024
0.354 0.398 0.550 0.624
0.1255 0.1586 0.3034 0.3901
From Emf Measurements 0.004462 0.0438 0.07673 0.01831 0.0679 0.1484 0.04174 0.0800 0.2157 0.07431 0.0814 0.2910
0.5708 0.4575 0.3709 0.2797
From Isopiestic Measurements 0.0786 0.0790 0.0618 0.0515
0.0908 0.1148 0.2156 0.2742
0.3102 0.3455 0.4721 0.5199
0.2534 0.2287 0.1309 0.0991
0.09 and 0.27 mCaCIZ. From these, values of (1 - R)/ ( Y N ~ c I * ~ Ncan ~ c ~be) ”calculated ~ (Table I). Gordon, et al.,596from emf and transference number measurements, have calculated osmotic coefficients for both sodium chloride and calcium chloride solutions. For example, Janz and Gordon give (b = 0.9414for a 0.0576 m solution of sodium chloride. By interpolation in the data of McLeod and Gordon, a solution of 0.04174 m calcium chloride has the same vapor pressure as this sodium chloride solution. In this way, values of (1 - R)/(YNaC1’mNaIJ1)1’2 were calculated for four pairs of solutions (Table I). Figure 1 is the plot of (1 - R ) / ~ / = R against l/m~ for both the isopiestic and the emf measurements 0.8
I
I
I
I
1
I
-
0.7 t \
\ \
0.60.5
c.3
0 2 0.1
0
From emf measurements
-
\\
1
4
t-
c
1
1
i
0
0
31
02
03
m
04
Figure 1. Plot of (1 - R ) / l / z us. ~ chloride solutions (R = NaCl).
The Journal of Physical chemistry
05
06
fi~ for calcium
07
of calcium chloride solutions. The extrapolation in Figure 1 shows that the integrand in eq 1 does approach -0.716. Suppose we did not have Gordon’s data for calcium chloride below 0.1 m and had only the points from Stokes’ work; t o know the limiting value is -0.716 is useful, but it does not completely solve the problem of filling in the curve between 0 and 0.1 m. In fact, there is much need for more work of Gordon’s type for extrapolations in the dilute region. Acknowledgment. The author wishes to thank Professor R. A. Robinson for suggesting improvements in the manuscript, and Professor C. W. McDonald for helpful discussions. (5) G. J. Janz and A. R. Gordon, J . Amer. Chem. Soc., 65,218 (1943). (6)
H. G. McLeod and A. R. Gordon, ibid., 68,58 (1946).
Identification of the Photochemically Active Species in Sulfur Dioxide Photolysis within the First Allowed Absorption Band by Sachiko Okuda, T. Navaneeth Rao, David H. Slater, and Jack G. Calvert Chemistry Department, The Ohio State University, Columbus, O h h (Received June 83,1069)
43810
Recently the photochemistry of sulfur dioxide has been studied extensively by several research g r o u p ~ . l - ~ The rate constants for the quenching reactions of the photoexcited singlet (‘SOZ)and triplet (aSO,) states of A small fraction sulfur dioxide have been of the SOrquenched molecules of sulfur dioxide excited or 2537 react to form sulfur trioxide’ a t 3130
As
(1) (a) H. D. Mettee, J . Chem. Phys., 49, 1784 (1968); (b) H. D. Mettee, J . Phys. Chem., 73, 1071 (1969). (2) 5. J. Strickler and D. B. Howell, J . Chem. Phys., 49, 1947 (1968). (3) (a) T. N. Rao, S. S. Collier, and J. G. Calvert, J . Amer. Chem. SOC.,91, 1609 (1969); (b) T. N. Rao, S. 8. Collier, and J. G. Calvert, ibid., 91, 1616 (1969). (4) J. N. Disco11 and P. Warneck, J . Phys. Chem., 72, 3736 (1968). (5) T. C. Hall, Jr., “Photochemical Studies of Nitrogen Dioxide and Sulfur Dioxide,” Ph.D. Thesis, Univ. Calif., Loa Angeles, 1953. (6) P. J. Warneck, G. C. A. Technology Division, Bedford, Mass., personal communication to one of the authors. (7) The present experimental evidence suggests but certainly does not require that the oxidation product of SOz formed photochemically is 503. All analytical systems which have been employed in the previous and present study depend on analysis for S04a- formed by reaction of the product with water, Direct spectroscopic evidence for SO8 formationin this system has never been reported to our knowledge; the fact that the ultraviolet and infrared spectra of 80s are uncertain and appear to contain no highly characteristic structure complicates its unambiguous identification. Conceivably, the primary product may be a metastable dimer of SOz (08-0-S-0) or even the reactive tetroxide of sulfur (SO*). Until further more definitive evidence of the nature of the SOZ photoproducts is available, we will assume that SOa2- arises from the simplest, traditionally proposed product, 808.
NOTES
4413
(@'SO, og0.04,5 0.0P). At wavelengths greater than 2180 A photodissociation of sulfur dioxide is energetically impossible, and it is probable that sulfur trioxide formation in this case results from an oxygen atom exchange reaction involving one of the electronically excited states of sulfur dioxide (SO2*)
In our previous work it was found that low pressures of biacetyl quenched efficiently the excited triplet sulfur dioxide molecules by energy transfer, reaction 9.
+
%02 Ac2 -+ SO2
+ 3A~2
(9)
Under the same conditions the concentration of singlet molecules was not affected d e t e ~ t a b l y . It ~ ~became evisOz* SO2 --3 so3 so dent that a series of chemical experiments utilizing pure It is surprising in view of the extensive kinetic data now SO2 and SO2-biacetyl mixture photolyses could define available on the reactions of excited states of sulfur the nature of the reactive excited state of sulfur dioxide. dioxide that the spin state of the excited molecule inWe have determined the rates of photochemical forvolved in reaction 10 cannot be determined. The work mation of sulfur trioxide in a flow system using pure reported here provides this information for the first SO2 (740 mm) and SO2 (737 mm)-biacetyl (3.3 mm) time. mixtures at 25". The data are summarized in Figure The previous studies are consistent with the following 1. There is a marked lowering of the rate of SO3 forreaction mechanism for SO2 photolysis within the first mation in runs with added biacetyl. This lowering is allowed absorption band not an artifact associated with the presence and reaction of the carbonyl compound with sulfur oxides, since so2 hv + 'SO2 (14 within the experimental error there is no effect on the 'SO2 so2 + (2502) (1b) rate of sulfur trioxide formation with added acetone (9.5 mm) in the sulfur dioxide (730 mm). In this case -+ 3S02 SO2 (2) of course the triplet energy (80.5 f 1.5 kcal/mol)8 is 'SO2 +so2 hvr (3) much higher than that of the sulfur dioxide (73.7 kcal/ 9 and triplet energy transfer from sulfur dioxide mole), +so2 (4) to acetone is expected to be very slow. The rate of -+ 3S02 (5) sulfur trioxide formation is lowered with biacetyl addition to about 65% of that with pure sulfur dioxide 3S02+SO2 hv, (6) photolysis. This result is in good agreement with that +so2 (7) expected if the 3S02is the only reactive precursor to sulfur trioxide formation. From the rate constants for as02 SO2 + (2502) (8) reactions 8 and 9 measured previously ( k 8 / k 9 = 2.7 X The indeterminant product (2S02) in reactions 1 and 8 10-3),3ait can be calculated that about 63y0 of the can be ground-state sulfur dioxide, SO3 and SO, and/or triplet SO2 molecules and less than 1% of the singlet possibly some other chemical products which do not remolecules would be quenched by biacetyl addition a t generate any excited state of sulfur dioxide.? Although the concentrations employed here. estimates of the rate constants for each of the reactions The present data lead to a somewhat higher estimate 1-8 have been made and quantum yields of sulfur triof the quantum yield of sulfur trioxide formation than oxide determined, it remains impossible to establish the that reported by Hall5 and Warneck6; from the data of nature of the precursor to sulfur trioxide formation from 0.02 for the band of waveFigure 1, @goa = 0.08 the existing data. This can be seen from the following This value is lengths employed here (3126-2537 considerations. The above mechanism and, recent data surprisingly close to our recent estimate of the quantu? from sulfur dioxide photolyses at 2875 A, 25", and yield of triplet sulfur dioxide formation at 2875 A; Psoz 1 200 p , suggest that the ratio of concentrations a 3 3 9 0 2 = 0.080 ct 0.014. The near equality of these two of triplet to excited singlet sulfur dioxide is given by3& quantities may be somewhat fortuitous since the quantum yield of 3S02may be wavelength dependent. [3S02]/['S02] G kz/ks 72 However, all of the present data do establish that the The rate constant data show that the rate of quench3S02molecule is the dominant chemically active species ing of 3S02is much slower than that for lS02 in irradiwhich leads to sulfur trioxide in sulfur dioxide photolated sulfur dioxide3& ysis within the first allowed absorption band. The rate of 3S02 quenching by SO2 -~ ['SOzIk8 o. 090 data further suggest that practically all of the triplets formed in the photolysis of pure sulfur dioxide at high rate of lS02 quenching by SO2 ['SOz]h pressures react to give sulfur trioxide product. Since the measured quantum yield of sulfur trioxide is less than 0.09 (0.04-0.05), the quenching rate data do (8) R. F. Borkman and D. R. Kearns, J. Chem. Phgs., 44, 945 (1966). not shed any light on the state or states responsible for (9) G.Heraberg, "Molecular Spectra and Molecular Structure. 111. the sulfur trioxide formation. Either one or both Electronic Spectra and Electronic Structure of Polyatomic Molecould contribute to it. cules," D. Van Nostrand Co., Princeton, N. J., 1966,p 606.
+
+
+ +
+
+
+
+
*
s).
~
'Volume 73, Number 13
December 1969
4414
NOTES
The present results, coupled with considerations of energy requirements and spin conservation, favor the following more detailed description of reaction 8a. 350~
+ soz
---it
so3+ s0(32-)
(84
With vibrationally relaxed triplet molecules, reaction 8a is exothermic by 26.1 kcal/mol. We are attempting to test the mechanism and the stoichiometry suggested by (sa) through kinetic spectroscopy of the presumed product SO using a doubled ruby laser to excite directly the triplet state of sulfur dioxide through absorption within a forbidden singlet-triplet band of SOz.l o It will be most interesting to determine the spin state responsible for sulfinic acid formation in sulfur dioxidehydrocarbon mixture photolysis. In this case the quantum yields of product climb from 0.006 for methane to 0.26 with pentane.ll Unless a chain reaction is involved here, the singlet excited state must be an important reactant in this case, since the quantum yield of triplet sulfur dioxide formation is only 0.08.3& We are planning to investigate this system in an analogous fashion to that used in this study. Experimental Section The photolyses were carried out in a cylindrical quartz cell (4.4 cm diameter, 5.2 cm length). The initiating light was generated from a medium-pressure mercury arc (Hanovia, Type A, 673, 500 W), filtered through a Corning filter 7-54 (9863). The relative intensities of the different wavelengths of light absorbed by the SOz were as follows: 3126-32, 39%; 3022-28, 18%; 2967, 12%; 2894,4%; 2804,6%; 2753,2%; 2700, 2%; 2652-5, 8%; 2571, 2%; and 2537 8, 6%. Practically all of the light at each of the wavelengths was absorbed by sulfur dioxide in the pure SOZ,SO2--MezCO, and S02-Ac2 mixture photolyses and almost none by biacetyl or acetone for these conditions. Tank sulfur dioxide (Matheson Co.) was dried, freed from dust, and then passed through the photolysis cell at a rate of 10 cc/min. The SOz excess and the fraction of SO3 product which remained gaseous flowed into two bubblers which were placed in series; each bubbler contained 10 cc of an aqueous solution (10% ethanol) which was 0.1 M in HgClz and 1 M in HC1. Following the run, nitrogen gas (20 cc/min) was used to flush the system for 20 min. The cell and associated glass tubing leading to the traps was disassembled after each run and the walls washed with small portions of the mercuric chloride solution. The stable complex formed between the mercuric chloride and sulfur dioxide retarded the thermal oxidation of SO2 to so3 during analysis. The trap solutions and the wash solutions from the cell and tubing were analyzed separately by adding 210 mg of BaC12.2H20 solid to each 10 cc of solution while it was magnetically stirred. The apparent optical density (usually 0.05-0.3) of the Bas04 suspension at 475 mp The Journal of Physical Chemistry
oV-'
I
IO
I
I
I
80
30
40
T i m e , min,
Figure 1. The time dependence of the product SO8 (measured as SOdZ-) which is formed photoch6mically (3126-2537 A) in pure SO2 a t 740 mm, circles; in mixtures of SO2 (730 mm) and acetone (9.5 mm), squares; and in mixtures of SO2 (737 m m ) and biacetyl (3.3 mm), triangles. Temperature, 25'; intensity of light absorhed by SO*, 2.9 X quanta/sec.
.
was measured on a spectrophotometer after a regulated period of time. Calibration of the analytical system was made using known amounts of soluble sulfate salt under identical conditions. The method gave excellent reproducibility on known solutions. l2 Scatter observed in the photochemical data is believed to be a consequence of the difficulties inherent in the quantitative transport of SO3 product to the analytical system. Usually the majority of the SO3 product (-85%) appeared in the first trap with the majority of the remaining product divided between the cell and the connecting tubing between the cell and traps. In runs with added carbonyl compound the SO2 gas was partially saturated with the desired vapor by leading the gas over a sample of either the pure biacetyl at 0" or acetone at -3.2". About 30% of saturation was achieved for these conditions; the composition of the mixtures from run to run was found to be very reproducible. Analysis of the biacetyl-SO2 mixtures was carried out by phosphorescence excitation using standard mixtures of Ac2 and SO2 and the Turner spectrophosphorimeter (Model 210). The acetone-SOz mixtures were analyzed chromatographically on a 10-ft Porapak-Q temperature-programmed column. In the runs with added biacetyl and acetone a larger part of the SO3was found in the solution prepared from the cell washings. Apparently the polar vapors accentuated the condensation of so3 in these cases. Following the analysis of a given run and before each (10) The absorption spectrum of SO has been observed in the flash photolysis of S0a by R. G. W. Norrish and G. A. Oldershaw, Proc. Roy. Soc., A249, 498 (1959); however, in this study the authors attribute SO formation largely to the direct photodissociation of SO2 at wavelengths shorter than 2200 A present in the flash and to the photolysis of some high energy isomer of SOa. (11) F. 8. Dainton and K. J. Ivin, Trans. Faraday SOC., 46, 374 (1950). (12) The analytical system was based in part on the work of H. J. Keily and L. B. Rogers, Anal. Chem., 27,759 (1955).
NOTES new run, the cell and connecting tubing were washed thoroughly, dried, reassembled, and evacuated overmm. night at a pressure of less than Dark runs were made which were identical in other respects with the photochemical runs. The amount of so3 derived from these experiments was considerably smaller than that formed photochemically and was very reproducible. For a dark run of given duration the amount of SO3 formed thermally was the same for pure SOZ, SO2-MeZCO, and SOz-Ac2 mixtures within the experiment,al error. Correction for SO3 formed thermally has been made in the data for the photochemical runs shown in Figure 1. Actinometry was carried out using the potassium ferrioxalate system.la Correction was made in the quantum yield calculations for the 3650-60 and 3341-A light present in the photolysis beam which was absorbed by the actinometer solutions but not by the SOz. The light intensity absorbed by sulfur dioxide was essentially constant throughout the series of runs: 2.9 X quanta/sec (3126-2537 A). Acknowledgment. The authors acknowledge gratefully the support of this work through a research grant from the National Air Pollution Control Administration, U. S. Department of Health, Education, and Welfare, Public Health Service, Arlington, Va. We are indebted to Professor Paul Urone (University of Colorado) for helpful suggestions which formed the basis for the development of the analytical system employed, and Dr. James W. Gall, who performed the SOz-acetone mixture analyses. (13) C. G. Hatchard and C. A. Parker, Proc. Roy. SOC.,A235, 518 (1956).
Anion Exchange of Metal Complexes.
X1X.l Volumetric Studies of the Exchanger in Mixed Solvents by Y. Marcus, Depurtment of Inorganic and Analytical Chemistry, The Hebrew University, Jerusalem, Israel
J. Naveh, and Mayo Nissim Israel Atomic Energy Commission, Nuclear Research Center-Negev, Israel (Received June $0,1060)
In a recent paper2 we reported the selective swelling of a divinylbenzene polystyrene-methylene-trimethylammonium salt copolymer (Dowex-1) in several aqueous organic solvent mixtures. This information is essential for an understanding of the factors affecting the sorption of metal complexes on the exchanger, which
4415 is of great practical importance. The applicability of anion exchange of metal complexes in mixed solvents depends, among other factors, on the rate of particle diffusion and on the dimensional stability of the column of resin, and these quantities depend, in turn, on the specific volume of the exchanger and on the solvent composition. It is further expected that some useful information concerning the interaction of the water and organic solvent with resin functional groups can be obtained from a study of the partial molar volumes of these components as a function of composition. These, again, are obtainable from the specific volumes. A study has therefore been undertaken of the density of the swollen resin, in equilibrium with a mixed solvent of known composition. The densities of 4 and 8% cross-linked resins in chloride form were measured in water and in aqueous methanol, ethanol, n-propyl alcohol, acetone, and formamide, a t 22’. Experimental Section Densities were measured in 25-ml picnometers at 22 f l o , the displaced liquid being the equilibrium solution itself, the density of which was separately measured. I n this way there is no fear from changes in the composition when the density of the swollen resin is measured. The effect of the temperature variations on the densities were within the reported experimental precision. The characterization of the resin and other materials, the procedure for drying the resin, and other operations have been reported.2 Calculations The specific volume, that is the volume per gram of dry chloride form of resin of a,resin sample swollen in a mixed solvent, is given by
P=
PR
+ nsPs’ + nwPw’
(1) where PR is the specific volume of the resin skeleton with its functional groups, assumed invariant with solvent composition, ng and nw are the specific numbers of moles of solvent (S) and water (W), and Vs‘ and Vw’ are the partial molar volumes, respectively. The specific swollen volume is obtained experimentally as
T‘
=
(1.000
+ nsMs + nwMw)/d
(2)
where M s and MW are the molecular weights and cl the measured density of the swollen resin. The molar volume of the solvent in the resin is now obtained from while a reference molar volume is calculated as
(1) Previous paper in series: Y. Marcus and E. Eyal, J . Znorg. Nucl. Chem., submitted, 1969. (2) Y. Marcus and J. Naveh, J . Phys. Chem., 73,591 (1969).
Volume 75, Number 12 December 1069
