Reaction Rates in the Analytical Determination of Some inorganic Peroxides and Superoxides SHELDON H. COHEN and JOHN L. MARGRAVE Deparfmenf o f Chemistry, Universify o f Wisconsin, Madison, Wis.
b The method for detection of small amounts of superoxide impurity in peroxides suggested b y George has been tested by observing the rates of oxygen evolution when lithium peroxide, sodium peroxide (both a white sample and a yellow commercial sample), sodium superoxide, potassium superoxide, strontium peroxide, and barium peroxide are dissolved in water. The resulting data show a ,definite difference in the rate of liberation of superoxide oxygen as dis9inguished from peroxide oxygen, but do not support the quantitative nature of the method as reported previously. It is unlikely that the sodium superoxide content of the yellow commercial sodium peroxide sample studied is greater than 6%.
T
and alkaline earth metals have been found to form either peroxides or superoxides, or both, on oxidation under high oxygen pressures. The thermodynamics of these materials has been investigated (4), but most studies of peroxides and superoxides have been beset by the difficulties of analysis. I n two recent papers George (2, 3) has described a method for determination of the percentage of superoxides in mixtures with peroxides by observation of the rate of reaction with water, especially immediately after mixing. For a sodium peroxide-sodium superoxide mixture the pertinent reactions are 2 Na02 2 H20 = H202 2 NaOH -I- 0 2 (1) N&O2 2 HzO = HZOZ 2 NaOH (2) HzOz = Hz0 '/z Oz (3) HE ALKALI
+ +
+
+
+
According t o George, the first two reactions are relatively fast but the decomposition of hydrogen peroxide is a t least 600 times slower unless a catalyst is present. Thus, solution of a peroxide should not produce oxygen, while a superoxide dissolved in water should liberate the superoxide oxygen rapidly, and a mixture should be analyzable. George has reported studies on a commercial yellow sodium peroxide, for which he suggested a superoxide content of about IO%, and a sample of potassium superoxide which presumably 1462
ANALYTICAL CHEMISTRY
contained 61% superoxide and 32.5% peroxide. Other workers have not clearly established the existence of potassium peroxide. Pure white sodium peroxide has been prepared a t the Callery Chemical Co. by atomizing liquid sodium into an oxygen atmosphere in the apparatus used for commercial production of potassium superoxide (7). Commercial yellow peroxides are prepared by stepwise oxidation to NazO, then Naz02, a t approximately 1 atm. EXPERIMENTAL
Apparatus. A modified gas buret calibrated in 0.05-ml. divisions and surrounded by a temperature-control jacket was connected to a glass reaction chamber by a glass tube fitted with a stopcock. The solid sample (0.1 to 0.5 gram) and water were held separately in small tubes and mixed by inverting the water container. I n all runs the solid peroxides and superoxides were loaded into the sample holder in a dry box, stoppered, and removed, then quickly attached to the system and made to react with the water. The volume of oxygen evolved was measured as a function of time and corrected to standard pressure and temperature. Materials. Samples of sodium superoxide, potassium superoxide, and
A number of commercial sodium peroxides were studied. The purity of such materials has always been subject to considerable uncertainty, since the x-ray pattern for pure sodium peroxide has only recently been established and analyzed ( I ) ; both sodium monoxide and sodium superoxide are likely impurities; and small amounts of sodium hydroxide in sodium peroxide are very hard to detect. In terms of active oxygen, a freshly opened can of sodium peroxide runs about 92 to 95% pure. Spot tests made on a number of samples indicated that no appreciable concentrations of iron or potassium were present. High purity samples of lithium peroxide, strontium peroxide, and barium peroxide were made available by Barium and Chemicals, Inc.
PURE
200-
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4a
(I)
B
2
white sodium peroxide were available from the Callery Chemical Co. The purity of the sodium superoxide was about 92% based on total oxygen evolution in mater in the presence of manganese dioxide. The potassium superoxide was about 89% *pure by similar analysis. The superoxides were strongly paramagnetic. The other impurities were carbonate, hydroxide, and possibly some hydrates. The white sodium peroxide was about 9401, pure and was diamagnetic.
150-
0
a
-+ W
n
n
I?
0 " -
k
100-
i TIME IN S E C O N D S
Figure 1.
Rate of oxygen evolution from superoxides
solution process. Magnetic studies also showed the lithium peroxide to be diamagnetic, White sodium peroxide reacted with water also yielded appreciable oxygen in a rapid process and leveled off a t 8.34 ml. per gram. If this \+eresuperoxide oxygen, it mould correspond to 4.1% superoxide. Seyb and Kleinberg (8)have studied this same white material by decomposing it rTith an acetic acid-diethyl phthalate mixture and found no superoxide oxygen; on reaction n-ith water they report evolution of 7.7 and 12.0 ml. per gram. The nhite material has been found t o be diamagnetic, indicating no appreciable paramagnetic superoxides or iron, although experimental error could allow up to 1% superoxide. Yellow sodium peroxide reacted with water to liberate more oxygen than the white sample, and calculation of the apparent sodium superoxide indicated about IO%, a value close to that reported by George for his samples. Yellow sodium peroxide samples appeared to be slightly paramagnetic.
DATA AND RESULTS
I n order to verify the suggestion of George (d, 3) that only the superoxide oxygen is evolved on solution of water, all the compounds listed above were made to react with water and the rates of evolution of oxygen measured. Strontium peroxide and barium peroxide gave no apparent gas evolution. The rate of oxygen evolution for sodium and potassiuni superoxides is plotted in Figure 1 as a function of time after mixing the solid with water. The oxygen evoiution by the superoxides is essentially complete after about 100 seconds. As only oxygen as predicted by Equation 1 is liberated rapidly and gas evolution then levels off, Reaction 3 is not proceeding at an appreciable rate. Therefore, no catalyst for hydrogen peroxide decomposition is present, and the time lag of the apparatus is unimportant in times of the order of 20 seconds or longer. The corresponding rate data for lithium peroxide, white sodium peroxide, and yellow commercial sodium peroxide are presented in Figure 2. I n the case of lithium peroxide some oxygen was rapidly evolved, apparently leveling off a t about 2.58 ml. per gram of sample. As no superoxide of lithium has ever been prep:rred in the solid form, i t is not possible to explain this oxygen except to assume that it comes from partial decomposition of the hydrogen peroxide formed, probably during the
DISCUSSION
From this work and the studies of Seyb and Kleinberg ( 8 ) ,it appears that George's (8) method for analysis of mixtures of peroxides and superoxides must be carefully considered before use on lithium or sodium compounds, since white, presumably superoxide-free per-
oxides of both metals still yield oxygen rapidly on reaction hith water. AS a first approximation, one may assume that the oxygen liberated by white sodium peroxide is from decomposition of the hydrogen peroxide formed, and place an upper limit of 6% on the sodium superoxide content of the yellow commercial sodium peroxide samples studied. Mechanical mixtures of white sodium peroxide and yellow sodium superoxide prepared for color comparisons also favor a 5% or lower superoxide content. I n addition, if sodium superoxide were present in commercial sodium peroxide to the extent of lo%, there should be easily detectable x-ray evidence. This has not been found, despite extensive workon the NazOz-h'aOz system (g), although a phase Na20p+. cannot be definitely excluded. Unfortunately, the best thermodynamic data available make it appear that sodium oxidized in low-pressure oxygen a t temperatures above 100" C. should not form any superoxide a t all, although the equilibrium data and high temperature heat capacities are only approximately known (6). The estimated equilibrium oxygen pressure over sodium superoxide a t 298" K. is 4 4 atm. and a t 600" K. rises to 40 =k 20 atm. Further high temperature equilibrium and heat capacity measurements along with phase studies of the type reported by Rode and Gol'der (6) will be required to explain sodium superoxide contamination of commercial sodium peroxide.
*
2c ACKNOWLEDGMENT
The authors wish to acknowledge the aid of William H. Schechter, Callery Chemical Co. and Andrew Pavlik, Jr., Barium and Chemicals, Inc., for supplying the samples of peroxides and superoxides used in this work.
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LITERATURE CITED
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Bailey, S. W., Tallman, R. L., Margrave, J. L., J . Am. Chem. SOC.
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Gc?orge, P., Discussions Faraday S O ~ . 2, 196 (1947). G€!orge, P., J . Chem. SOC.70, 2367 (1955). Gilles, P. W., Margrave, J L., J . Phys. Chem. 60. 1333 (1956). Margrave, J . L., J . Chem. Educ. 32, 522 (1955). Rode, T. V., Gol'der, G. A., Izvest. Akad. Nauk. SS.S.R., Otdel. Khim. Nauk 3, 299 (1956). Schechter, W. H., private communication, 1956. Seyb, E., Kleinberg, J., ANAL.CHEM. 23, 115 (1951). Tallman, R. L., unpublished work, University of Wisconsin, 1954-57.
TIME IN SECONDS
Figure 2.
Rate of oxygen evolution from peroxides
RECEIVED for review November 23, 1956. Accepted May 8, 1957. VOL. 29, NO. 10, OCTOBER 1957
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