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Water solutions containing ceric perchlorate and perchloric acid evolve oxygen when they absorb ultraviolet light and the ceric is reduced to cerous perchlorate.2Baur2a and Weiss and Porret2b studied the photochemical reaction in the full light of a quartz mercury arc lamp. Weiss and Porret obtained maximum gross quantum yields of the order of one-tenth in solutions one-tenth molar in ceric perchlorate and one molar in perchloric acid. Their quantum yields decreased as the cerous perchlorate accumulated.We have measured the quantum yields of the reaction when the solutions are irradiated with monochromatic light of X 254 µ. The light intensity, I, and the concentrations of cerous, c3, and ceric, Ci, perchlorates were varied many fold. The perchloric acid concentration, c2, was held at 1.03 =*= 0.03 M and the ionic strength, µ, at 1.1 ± 0.1, all at 23 ± 3°.Solutions of ceric perchlorate in perchloric acid are thermally unstable at 25°. The equilibrium ratio Ci/Ci is about 10 _8 in molar perchloric acid in equilibrium with the atmosphere (po, = 0.2 atm.), but the rate of the thermal reduction of the ceric perchlorate by water is extremely slow. In one of our stock solutions which was 5.3 M in perchloric acid and was kept in the dark at 25 * 3°, Ci decreased from 1.36 to 1.20 M in eleven months while c3 increased from 0.32 M.Materials.-The chemical reagents were of analytical reagent grade or were prepared from material of this quality. The water was chloride-free distilled water.Stock solutions of ceric and cerous perchlorates were prepared from a sample of snow white granular ceric oxide, 98.5% pure Ce02, which was supplied by the Rohm and Haas Chemical Co., Philadelphia, Pa. The ceric oxide could not be converted directly into ceric perchlorate even when dispersed as a fine hydrous oxide in 72% perchloric acid at room temperature or at 100°f or periods of several months.The ceric perchlorate was finally prepared by reducing the ceric oxide to cerous ions by bromide in perchloric acid. The bromine and excess bromide were removed by boiling. The cerous ions were then oxidized electrolytically to the ceric state. The experimental details follow.Forty grams of the granular ceric oxide and 80 g. of sodium bromide were added to 275 ml. of 72% perchloric acid. The mixture was simmered for two hours under an appropriate hood. The hot solution was filtered by suction through a sintered glass filter. The filtrate was boiled(1) The part of this article concerned with the experimental study of the effect of cerous perchlorate upon the reaction is taken from the thesis submitted by Maynard E. Smith in September, 1946, to the Department of Chemistry of the Massachusetts Institute of Technology in partial fulfillment of the requirements for the degree of Master of Science.(2) (a) E. Baur, Z. physik. Chetn., 63, 683 (1908), was the first to identify oxygen as the gaseous product of the reaction. This was

A study has been made at 25" of the influence of pH upon the relative quantum yields for the consumption of oxalate and the production of carbon dioxide, carbon monoxide, and uranous ion by light of 254 nm absorbed by the uranyl oxalate actinometer system. The uranyl ion was at 0.01 F , the oxalate at 0.06 F , and initial pH was at 0-6. About 12% of the oxalate was decomposed. The quantum yields for the consumption of oxalate were found to be independent of the pH between 1 and 5 but to decrease outside this range. The reaction at all values of pH was found to consume acid. The moles of carbon dioxide produced in all forms (COZ, H&Oa, HCOa' , COa2-and the uranyl carbonate complexes) per mole of oxalate consumed very nearly equalled unity at pH 0 to 5. The moles of carbon monoxide produced per mole of oxalate consumed were always less than unity; they decreased with increase in pH abruptly between pH 1 and 2, and became negligible above pH 3. The mole ratio of uranous ion produced to oxalate consumed increased slowly from about 0.03 between pH 1.5 and 4 to about 0.08 at pH 6. The mechanisms of the reactions are discussed.

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Continued experience with the uranyl oxalate actinometer has led us to vary the concentration of its components, the better to adapt them to the frequency and intensity of light employed. We have measured the quantum yields, 4, of such solutions in terms of d,s for a "standard" solution, 0.01 M in uranyl sulfate and 0.05 M in oxalic acid, carefully investigated. 1*2 Commercial uranyl sulfate is often of dubious quality, and then requires laborious purification.2 Heidt and Daniels2 obtained uranyl oxalate in sufficient purity merely by combining commercial uranyl nitrate with oxalic acid, washing the precipitate and drying at 110'. They photolyzed a solution 0.01 111 in uranyl oxalate and 0.05 114 in oxalic acid at 1313 mp, and found I$ = 0.57, identical with I#I~ found by them and by Leighton and Forbes1 a t the same wave length using uranyl sulfate.There are at least three good reasons for introducing less than 0.05 mole of oxalic acid per liter.(1) The number of quanta absorbed by the actinometer must be calculated from a relatively small difference between two titrations of total oxalate. If this total is decreased, a smaller difference can be determined with equal percentage accuracy, and the time required for actinometry in the course of a photochemical experiment can be shortened.( 2 ) Excess of oxalate exerts a strong inner filter effect if X < 254 mp. Thus at A208 mp the extinction coefficient K H~c~o , = 1300 as compared with KUO,C~O~ = 4500. (3) at A208 mp unsensitized photolysis of oxalic acid occurs (d, = 0.02).3 For reasons (2) and (3) confusion results if concentrations vary.If less than five molecules of the weak oxalic acid are added to one of uranyl sulfate, liberation of hydrogen ion brings the reaction to equilibrium short of "complete" conversion to uranyl oxalate. Consequently the inner filter effect of uncombined uranyl ion is not at a minimum, nor the concentration of uranyl oxalate (presumably the photolyte) at a maximum. Then d, falls below d,, and becomes :t function of ~oxicentration.~ But if one starts (I) W.G.Leightonand G.S Forbes,Tars JOURNAL,^^,^^^^ (1930) (2) Heidt and Daniels, ibid ,, 54, 2384 (1932).(3) Brackett and Forbes, ibid., 56, 4459 (1933) (4) Bucbi, 2 physik Chcm , 111, 269 (1924) with 0.005 M uranyl oxalate, and maintains even half a molecule of oxalic acid in excess throughout the photolysis, $I can be assumed constant at d,, as shown below. The advantages of restricting excess of oxalate can thus be realized without encountering diminished and variable quantum yields.No great decrease in oxalate concentration is feasible through cutting down uranyl concentration when 436 mp > X > 366 mp because transmissions become inconveniently large unless reaction layers are made unduly thick. But when X = 205 mp, "standard" solution absorbs 7770 of incident light 0.01 mm. from the front window. Vigorous stirring is then necessary to avoid depletion of oxalate, liberation of uranyl ion and consequent decrease of d, in this thin layer. Using 0.001 ,I1 uranyl sulfate or oxa...