and
+ :w+ ki JI k2 R 0+ Y
I W Z
The main requirement for a time dependent phase angle with such systems is that the d.c. process is under the influence of charge trrtnsfer or chemical reaction kinetics. T me dependence of phase angles may prove a useful diagnostic tool for such systems. Adsorption and second or higher order chemical kinetic effects also are suspected as possible sources of phase angle time dependence although presently only qualitative argument:; can be given. RESllLTS
We have undertaken an investigation of the validity of Matsuda’s theory regarding the time dependence of the a.c. wave. Preliminary experimental results support, a t least qualitatively, the theoretical conclusions. Experimental appa1 atus and circuitry employed in this work have been described elsewhere (14, 1 6 ) . All solutions were prepared from lrtboratory distilled water, which was redistilled from alkaline permanganate, and reagent grade compounds with one exception. The V(II1) solutions were prepared from vanadium trichloride crystal purchased from K and K Laboratories, Inc. Handling of vanadiuin trichloride and preparation of V(II1) solutions were performed in a dry box under a nitrogen atmosphere. Experiments with the ferric-ferrous system in 0.5M sodium oxalate [kh> 1 cm. second-’ ( 1 1 ) ] and the Ti(IV)/Ti
(111) system in 0.20OM oxalic acid [kA = 4.6 x lo-* cm. second-’ ( I C ) ] show no dependence of alternating current amplitude on mercury column height as would be expected because the heterogeneous charge transfer rate constants exceed 10-2 cm. second-’. However, the V(III)/V(II) system in 1M H2SO4 [ k ~= 1 X cm. second-’ ( l o ) ] shows a definite dependence on mercury column height which varies with d.c. potential as predicted by theory. Some typical results are shown in Figure 2. The cross-over potential, Ed.c, *, is independent of alternating potential frequency in agreement with theory. The value of a calculated from the Ed.?.*potential was 0.46 + 0.03, in excellent agreement with the value reported by Randles (10) for this system based on conventional faradaic impedance and d.c. polarographic measurements. These initial results show promise regarding the validity of the theory and the applicability of a x . polarographic time-dependence studies. Additional, more detailed, experimental investigations will be attempted. Among them will be observations on systems involving chemical kinetic and adsorption effects, quantitative examination of faradaic alternating current-time curves at the dropping mercury electrode, the dependence of phase angles on mercury column height, etc., and similar studies with higher ha,rmonic currents. LITERATURE CITED
(1) Berzins, T., Delahay, P., Z. Elektrochem. 59, 792 (1955). (2) Breyer, B., Bauer, H. H., “Chemical
Analysis,” Vol. 13, P. J. Elving and I. M. Kolthoff, eds., Chap. 2,. Inter. science, New York, 1963. (3) R. P.. J . Electroanal. Chem. . ,5,Buck. 295 11963). ’ (4) Delahay, P., in “Advances in Electrochemistry,, and Electrochemical Engineermg, Vol. I , P. Delahay and C. W. Tobias, eds., Chap. 5, Interscience, New York, 1961. (5) Hung, H. L., Delmastro, J. R., Smith, D. E., J . Eleclroanal. Chem., in press. (6) Kuta, J., Smoler, I., in “Advances in Polarography,” Vol. 1, I. S. Longmuir, ed., pp. 350-8, Pergamon, New York, 1960.
(7) Kuta, J., Smoler, I., in “Progress in Polarography,” Vol. 1, P. Zuman, ed., with collaboration of I. M. Kolthoff, Chap. 3, Interscience, New York, 1962. (8) Matsuda, H., 2. Elektrochem. 62, 977 (1958). (9) Matsuda, H., Ayabe, Y., Bull. Chem. SOC. J a p a n 28, 422 (1955). (10) Randles, J. E. B., Canadian J . Chem. 37, 238 (1959). (11) Randles, J. E. B., Somerton, K. W., Trans. Faraday soc. 48, 937 (1952). (12) Reilley, C. N., Stumy, W., in
“Progress in Polarography, Vol. 1, P. Zuman, ed., with collaboration of I. M. Kolthoff, Chap. 3, Interscience, New
York, 1962. (13) Smith, D. E., ANAL. CHEM.35, 602 (1963). (141 lbid.. D. 610. (15) Zbid.; ‘p. 1811. (16) Smith, D. E., Northwestern Uni-
versity, Evanston, Ill., unpublished work,-1963.
HOYINQ L. HUNQ DONALD E. SMITH
Department of Chemistry Northwestern University Evanston, Ill. RECEIVED for review December 20, 1963. Accepted February 3, 1964. Research supported by the National Science Foundation.
Effect of Metal tons on Nuclear Hyperfine Coupling Constants in EPR Spectra SIR:We have investigated the effects of metal ions on the nuclear hyperfine coupling constants of the electron paramagnetic resonance (.EPR) spectra of a variety of anions of aromatic nitro compounds. The anion radical of p-chloronitrobenzene was selected as a model system since it3 EPR spectrum consists of three well resolved triplets resulting from interaction of the unpaired electron with i,he N14 nucleus, two equivalent (ortho) protons and two more equivalent (meta) protons. The radical anions were generated electrochemically in dinethylformamide (DMF) containing the metal ions to be investigated as suppor ,ing electrolytes. The effect of the metal ion on the W 4 coupling constant (aN)was investigated.
Coupling constants are believed to have better than 1% relative error. The variation of U N is shown in Figure 1 as a function of concentration of the various supporting electrolytes. The figure shows first that the shifts in uN caused by metal ions are considerably greater than those occasioned by interaction with water or other hydrogen-bonding solvents when compared on a concentration basis (4). The shift in a N qualitatively follows the solvation energy and other physical parameters of the various cations in DMF. No metal ion couplings are observed in the E P R spectra and.bhe exact nature of the interaction between the anion radical and the metal cation species is
not known. A rapidly exchanging metal ion system should show a lack of metal ion coupling. Also, since all solutions contain about millimolar water under real experimental conditions, the interaction may be via aquo-metal ion species. The E P R spectra also show a line broadening effect most markedly observed in the odd spin states of the nitrogen. A similar line broadening does not occur in aqueous-dimethylformamide mixtures with tetraethylammonium perchlorate as supporting electrolyte, indicating that this effect is definitely caused by the presence of the metal ions. Metal ion effects are also noted in the cyclic voltammetry of the nitro compounds. The first peak, corresponding VOL. 36, NO. 4, APRIL 1964
925
CONC.
(NOLES/L.)
Figure 1. Variation of NI4 coupling constant with various metal ions
to the reversible addition of one electron to form the anion radical, is only vary slightly shifted toward more anodic potentials. However, the second peak is shifted markedly t o more anodic
potentials by, say, lithium ion. These results are completely analogous to those recently given by Peover and Davies (3)for the reduction of quinones in DMF and by Holleck and Becher for nitro compounds (2). These effects are important because, in reporting coupling constants of radical ions generated electrolytically, it will now be necessary t o include the supporting electrolyte if this is other than a tetraalkyl salt. The almost complete lack of U N dependence on tetraethylammonium perchlorate concentration (Figure 1) shows that the original choice of tetraalkyl ammonium salts as supporting electrolytes in electrolytic generation of radical anions by Geske and Maki was indeed a judicious one (1). Bowever, as a consequence of the metal ion effects on coupling constants, one is able to study metal ion-radical ion interactions via the EPR method which gives additional information on the physical picture of
the interaction. A detailed report of these studies will be given soon. LITERATURE CITED
(1) Geske, D. H., Maki, A. H., J . Am. Chem. SOC. 8 2 , 2671 (1960). (2) Holleck, L., Becher, D., J. Electroanal. Chem. 4, 321 (1962). (3) Peover, M. E., Davies, J. D., Zbid., 6, 46 (1963). (4) Piette, L. H., Ludwig, P., Adams, R. N., J . Am. Chem. SOC. 84, 4212 (1962).
TOYOKICHI KITAGAWA~ THOMAS LAYLOFF RALPHN. ADAMS Department of Chemistry The University of Kansas Lawrence, Kan. Present address, Department of Chemistry, Osaka City Univereity, Osaka, Japan.
RECEIVED for review January 10, 1964. Accepted January 23, 1964. This work supported by the Air Force through Air Force Office of Scientific Research and by the Atomic Energy Commission through contact AT( 11-1)686.
The Dissolution of Calcium Tungstate with an Acid-Hydrogen Peroxide Mixture SIR: The use of single crystal Ca-
wo4 as a laser host and the associated
problems of doping and charge compensation have created a demand for analytical methods for this material. Although a Na2C03-K2C03 fusion will render CaW04 soluble ( d ) , this procedure is time consuming and not applicable for low-level determinations of alkali metals as in the case with charge-compensated, rare earth-doped crystals. Consequently, a method for the rapid quantitative dissolution of CaW04, involving mild heating with strong acid and hydrogen peroxide, has been developed. This communication reports the details of this procedure, its mechanism, and suggests some possible extensions t o other materials. EXPERIMENTAL
Two milliliters of concentrated acid (HC104 or HN03), 20 ml. of 3001, H202, and 50 ml. of distilled water are normally used for each gram of CaW04. However. the amounts of acid and peroxide ' used are not critical. The rate of dissolution depends critically on solution temperature and surface area of the sample. The optimum solution temperature is SO" to 90" C.-i.e., the highest teniperature which does not cause excessive decomposition of the hydrogen peroxide. Finely divided powder can be dissolved in 2 to 5 minutes while single crystals and sintered material ground with a mortar
926
0
ANALYTICAL CHEMISTRY
pestle require 15 to 30 minutes depending on the degree of fineness achieved. Solutions prepared in this fashion using HClOn are stable for up to 4 months depending on the amount of excess peroxide added. Slow continuous decomposition of the peroxide is evidenced by the slight effervescence of the solutions. Evcntually a yellow tungstic acid precipitate forms. Similar behavior is noted when the solution is boiled. The precipitate can be redissolved by cooling the solution and adding more peroxide. DISCUSSION
For applications such as flame photometry, samples and standards prepared in this manner can be used directly. However, in applications where the solutions must be heated either a t high temperature or for long periods of time, or in cases where the bubbling may be a problem-e.g., x-ray fluorescence spectrometryfurther treatment of the solutions is necessary to stabilize them. A simple method for accomplishing this is to add a slight excess of (ethylenedinitril0)tetraacetic acid (EDTA) t o the solution and then bring the p H of the solution to 12 by addition of base. The majority of the peroxide decomposes spontaneously in basic solution and can be removed entirely with mild heating. The dissolution appears to be a two-
step process. The CaWOa 1s decomposed forming tungstic acid which reacts with the peroxide to form the soluble peroxytungstic acid. No tungstic acid precipitate is formed during the dissolution. Addition of EDTA before making the solution basic complexes the Ca+?which would otherwise precipitate as Ca(OH)?in basic solution. As noted above, particle size and temperature are the most important parameters in the speed of the reaction. Since the primary reaction is the acid decomposition of CaW04, the greatest possible surface area should be presented for the reaction. An optimum temperature exists for the reaction because the CaWO4decomposition is more rapid a t high temperatures while excess heating destroys the hydrogen peroxide. Although stirring is desirable to promote the reaction, Teflon magnetic stirring bars should not be used because they are attacked by the reagent and/or abraded by the CaWO4. Small flakes of Teflon are inevitably found in the solution and repeated use of the same stirrer causes iron from the bar to be released into the solution. Three acids, hydrochloric, nitric, and perchloric, were tested for use in this method. Although the dissolution proceeded slightly faster with HC1 than with the others, effervescence of the resulting solution was more pronounced with this acid. Although no apparent
