3662
NOTES
atoms7 and that it occurs with zero activation energy above 800" (see Table I1 of ref. 1). Recently, Rosner and Allendorf have shown that the oxidation rate of molybdenum* and tungstens is also enhanced by 0 atoms. Their results are very similar to ours, and one would expect that the enhanced oxidation of these metals could be explained by a similar mechanism. There is some difference that results from the difference in stability of the oxides. The PtO2 is unstable a t the temperature of formation and must volatilize in a time comparable to the lifetime of the activated state or decompose. The probability of volatilizing in this short period is small so that the €0 for platinum is small, 5 X loF6. On the other hand, the oxides of tungsten and molybdenum are stable a t these temperatures and can exist on the metal surface for some time before volatilizing so that the primary reaction may be between adsorbed 3402 and a gas phase 0 atom. This reaction would result in the formation of the trioxide, which is compatible with the fact that metals are usually oxidized to their highest valence state in activated oxygen.2 (7) J. C. Greaves and J. W. Linnett, Trans. Faraday SOC.,54, 1323 (1958). (8) D. E. Rosner and H. D. Allendorf, J . Chem. Phys., 40, 3441 (1964). (9) D. E. Rosner, private communication.
The Oxygen Electrode in Fused Alkali Nitrates
by R. N. Kust' Department of Chemistry, Texas A & M University, College Station, Texas (Received M a y 3, 1966)
Several workers have reported on the existence of oxygen gas electrodes reversible to the oxide ion in fused s a l t ~ . ~ Generally, 13 the solvent used was one containing an oxyanion. However, accurate determinations of standard e.m.f.'s have not been made. Usually, the oxygen electrodes have been used in concentration type cells where the EO'S cancel out. The present author has published a study of the oxygen electrode in fused alkali nitrate solvents4 which included the E" for the oxygen electrode against a silver-silver ion glass membrane electrode. The cell reaction could be written 2Ag+
+ 0'-
2Ag
+
'/202
(1)
The E" was established by the coulometric addition Of oxide ion. Since the quantities of oxide ion generated The Journal of Physical Chemistry
were quite small, the oxide ion concentration ranging m, chemical methods of analyfrom about to sis were not employed to determine the oxide ion concentration. The amount of oxide ion added was determined by integration of the time over which a constant current of 10-20 pa. was passed through the cell. It was tacitly assumed that the current efficiency was 100%. It has been suggested that the definition of the electrode system would be more reliable if the values for the E" obtained coulometrically could be supported by data obtained from a chemical addition of oxide ion. This note is a report of such data. Experimental Section All chemicals used were of reagent grade. The solvent of equimolar sodium-potassium nitrates was prepared by fusing the proper proportions of the two salts and mixing well. Dry nitrogen was bubbled through the melt for 2 hr. The solvent was then filtered through a fine grade fritted glass disk, molded into slugs of about 100 g., and stored over magnesium perchlorate. The sodium carbonate used was dried at 300" for 4 hr. The reaction vessel and electrode were similar to those previously de~cribed.~An oxygen-platinum electrode was used as the indicator electrode and a silver-silver nitrate (1.0 m AgNO3 in equimolar Ya, KN03) electrode was used as the reference electrode. Potential measurements were made with a Leeds and Xorthrup K-3 universal potentiometer. A Keithley Model 603 electrometer amplifier was connected in series with the electrochemical cell and was used as a null detector. Results Small quantities of Na2C03were added to the electrochemical cell. The oxygen gas passing over the oxygen electrode swept out the COz generated by the dissociation of the carbonate ion according to the reaction co32-
co2 + 02-
(2)
Since oxide ion was the other product of the dissociation, the potential increased with time. The complete dissociation of the carbonate ion was indicated by the halt in the potential rise. The concentration of the oxide ion, calculated from the amount of sodium carbonate added, ranged from to m. Potentiometric readings were made after the e.m.f. had been (1) Department of Chemistry, University of Utah, Salt Lake City, Utah 84112. (2) H. Flood and T. Forland, Acta Chem. S c a d . , 1, 92 (1947). (3) B. A. Rose, G. J. Davis, and H. J. T. Ellingham, Discussions Faraday sot., 4, 154 (1948). (4) R. N. Kust and F. R. Duke, J. Am. Chem. SOC., 85,3338 (1963).
NOTES
3663
constant for a t least 20 min. The values of E" for the cell reaction were obtained at several temperatures and are listed in Table I. Each value for the E" is the average of at least five and in Some cases six measure ments of the e.m.f. at known oxide ion concentrations.
E o (v.)
= 0.7759
- 2.557
X lO-*T
530OK.
< T < 639%
With these additional data the oxygen electrode is Sufficiently defined to be a Useful electrode for electrochemical studies in alkali nitrate solvents. (6)R. N. Kust, Inorg. Chem., 3 , 1036 (1964).
Table I: Comparison of the Eo Values for the Oxygen Electrode Obtained by Chemical and Coulometric Methods EO,
v.
Temp., OK.
(chemical)
536 543.5 545 553 565 578 585 589 595 602 610 616 621 630 639
0.6390 0.6370 0.6367 0.6346 0.6314 0.6280 0.6260 0.6254 0.6236 0.6221 0.6200 0.6182 0.6171 0.6151 0.6128
Ea, V. (ooulometric)
0.6388 0.6369 0.6365 0.6345 0.6314 0.6281 0.6263 0.6253 0.6238 0.6220 0.6200 0.6185 0.6173 0.6150 0.6127
Discussion The chemical addition of oxide ion to an equimolar sodium-potassium nitrate melt is complicated by the insolubility of most metallic oxides. Also, the introduction of cations different from the solvent cations would lead to possible complex formation, the formation constants of which would be unknown. Hence, one is limited to either NaO or KzO as a source of oxide ion. It is very diflicult to prepare either of these oxides so that they are free from peroxide and superoxide contaminants. However, it has been shown that sodium carbonate has an unusually large dissociation constant in this solvent in the temperature range of i n t e r e ~ t . ~ The dissociation constant for reaction 2 is on the order of at 300". Thus the addition of Na&O, to the solvent and the subsequent removal of the COZ produced is equivalent to the addition of NazO directly. Also, no contamination by peroxides or superoxides is likely to occur. A comparison of the E" values obtained in this manner to the values obtained by the coulometric generation of oxide ion is given in Table I. I n every case the difference between the two values is 0.3 mv. or less. The E o can be calculated for any temperature between 530 and 639°K. from the equation
Aluminum-27 Nuclear Magnetic Resonance of Trialkylaluminum Compounds. 11. Variable-Temperature Studies
by Charles P. Poole, Jr., Harold E. Swift, and John F. Iteel, Jr. Gulf Reeearch & Developntent Company, Pittsburgh, Pennsylvania (Received May 10, 1966)
I n a previous publication' several aluminum alkyl compounds were studied by aluminum-27 n.m.r. both in the pure state and dissolved in various solvents. I n low-viscosity solvents the line width was found to be proportional to the viscosity times the cube of the molecular radius, and the dominant relaxation mechanism in these solvents was attributed to quadrupolar relaxation through molecular rotation. In high-viscosity solvents the line width became much less dependent on the viscosity. The present study employed variable-temperature techniques to obtain the temperature dependence of the line width of pure aluminum alkyls and mixtures of triethylaluminum in solution.
Experimental Section The n.m.r. measurements were made on a Varian V-4200-A wide-line n.m.r. spectrometer equipped with a V-4257 variable-temperature accessory. The experimental arrangement and spectrometer settings were identical with those employed in the room-temperature studies of these same chemical systems.' The temperature was monitored by a thermocouple located below the sample, and a correction was made for the temperature difference between the thermocouple position and the actual sample location. The sample tubes used in the variable-temperature studies had inside diameters of 8 mm., whereas the tubes used for the room-temperature studies had inside (1) C. P.Poole, Jr., H.E.Swift, and J. I?. Itael, Jr., J . Chem. Phys., 42,2676 (1966).
Volume 69, Number 10 Odober 1966
