Titrations were attempted with mixtures of halide ions. I n Figure 2 is bhown the titration curves of approximately cquul normnl i~oncrntrationiof ioditlr nntl bromide ion, atlmi\ed nith ortlioiilii~-l)hate 7’he prpciiioii of the hnliilp cnil puiiiti \\a\ not :ippre(+ildy Y I llin\rLer, t l i c b ort1iol)lic)sl’li:Lte end point5 occurred 23 and GO%, too soon, reyiectively, indicating con4 e r a b l c coprecipitation of the orthophohphate nith the silrer halide. Othrr poqsible interf 3rences 11ere also in\ c-tigated potentiorietricdly. Five milliliterb of 0.1.11 ortliophosphate mas titrated with 0 1JI s i l ~er nitrate in the preqence of an equal niolar quantity of the subdance being tested. The presence of sulfate lielded titration values : h u t 1 % greater than in it, absence. Khile this i. noticenbli., it appears that u p to equal molar q113 ntities of sulfate can be tolwatccl nithout great 10‘s of nccuraq . C‘alcium and aluminum each interfered niarkctllI\-. In t h ? case of calcium, the addition of a tno-fold excess of EDTd only partially eliminated the interference. Ferric iron also markedly affected the titration. The addition of a ten-fold excess of sodium fluoride again only partiallj eliminated the interference. Bniperonietric end point detection wab a1.o investigated. In this case, it n a b nece.sary to add gelatin to prevent depolarization of the ir dicator electrode by the sill cr phosphate precipitate. Silver (ahloride is known to depolarize a plntinuin electrode. Laitinen and co-
workers (7, 8 ) wed 0.1% gelatin to prevent depolarization by silver chloride. I n the titration of orthophosphate with silver ion, the best amperometric end points mere obtained in the presence of 0.2yogelatin. Even in the presence of gelatin, huwever, the silver phosphate precipitate depolarized the electrode somewhat, especially in the more dilute solutions. This is apparent in the amperometric titration curves to be discussed. Figure 3 shows typical amperometric and titration curve3 of 8.4 X 1.7 X 10-3X orthophosphate. The curve in 3 A was obtained by titration with 0.102M silver nitrate solution. The current readings were corrected for volume changes. Curve 3 B was obtained by coulometric generation of silver ion at a generation rate of 19.3 ma. The appreciable current readings noted as soon as qilver ion was added indicated some depolarization of the electrode by the silver phosphate formed. Because of this appreciable depolarization, a special procedure had to be used in obtaining the end point. Instead of extrapolating the current to zero and finding the end point a t this value, the base line had to be drawn through the line of points immediately preceding the end point, as in Figure 3 A . I n the case of the more dilute solutions of orthophosphate, there was noted a rounding of the current increase beyond the elid point. Therefore, the current had to be extrapolated in a curJed fashion, as shown in Fibure 3 b. This method was used with several titrations
and vias found to be reproducible. There was always an apparent current break between one third and one half of the end point value (see Figure 3 B ) . At phobphate concentrations less than 1.7 X 10-3.11, the amperometric end point was inadequate. I n a Yeries of coulometric and volumetric. titrations between the orthophosphate concentrations of 1.7 X 10-3 and 8.4 X l O - 3 J I , a relative error of 1% wa> obtained. LITERATURE CITED
(1) Brunisholz, G., Heln. Chim. Acta 30, 2028 (1947). ( 2 ) Calamari, J. .4.,Hubata, R., IXD. ESG.CHEAT., BSAL.ED. 14,55 (1942). (3) Cullum, D. C., Thomas, D. B., Anal. Chzm. ilcta 24. 205 11961). ( 4 ) Flatt, It., hrunishcilz,’ G., Ibitl., 1, 124 (1947). ( 5 ) Gerber, A. B., Miles, F. T., I r u . ENG. CHEM.,ANAL.ED.13,406 (1941). (6) Harrison, T. S., Parratt, T., J . SOC. Chem. Ind. iI,ondon) 64.218 (1935).
( 7 ) Laitinen, ‘ H. A.,’ Jennings, k. P., Parks, T. D., IND.ESG. CHEM.,AUL. ED. 18, 355 (1946). (8) Laitinen, H. A., Kolthoff, I. M., J . Phys. Chem. 45, 1079 (1941). (9) Moerk, F, Hughes, E., Am. J . Pharm. 94, 6 ~ (1922). 0 (10) Rancke-Madsen, E., Kjaergard, T., Acta Chem. Scand. 7 , 735 (19,53). (11) Sanfourche, iz., Focet, B., Bull. SOC. Chim. France 53,963 (1933). (12) Simmick, H., 2. dngew. Chem. 48, 566 (1935). (13) Thistlethwaith, W. P., Analyst 77, 48 11952). (14) Toller; W.,Mztt. Ver. Grosskesselbesztzer 1946, 88. (15) Wilson, H. X., Analyst 76, 65 (1951).
RECEIVEDfor review May 24, 1963. Accepted August 9, 1963.
Polarography of Nitroferrocene A. M. HARTLEY and R. E. VISCO1 Department of Chemistry and Chemical Engineering, University o f Illinois, Urbana, 111.
b The polarographic reduction of nitroferrocene in newiral or alkaline buffers is found to b e a diffusion controlled, concentraiion dependent, six-electron process. The half wave potential of the reduction is observed to shift 58 mv. per p H unit. In situ electrochemical reduction of nitroferrocene within an electroil spin resonance (e.s.r.1 spectrometer produces an unstable radical with a nitrogen coupling constant of 15.1 gauss. In aqueous buffers, nitroferrocene undergoes a photuchemical decomposition with 1 nitrocyclopentadieneide as one product.
-
s
THE DISCOVERY of the socalled L‘sandwich” compound, dicyclopentadienyl iron or “ferrocene,” numeroils investigations have been
INCC
concerned with the aromatic character of the parent compound and its derivatives. These studies have been reviewed by several workers (14, 18, 19). Recently, Rausch, Fischer, and Grubert (17) have compared the reactivities of ferrocene to the corresponding ruthenium and osmium compounds. It has been s h o m that ferrocene can undergo many of the reactions typical of aromatic benzenoid compounds such as Friedel-Crafts acylation and alkylation, sulfonation, arylation, and metallation with butyllithium. On the other hand, the compound cannot be considered completely a n aromatic substance since even the mildest of nitration conditions usually employed with aromatic compounds leads to destruction of ferrocene. Nitroferrocene has been prepared by Helling and Shechter (6) and by Grubert and Rinehart (4)using a method
which is unusual for aromatic nitrations. Ferrocenyllithium was treated with n-propyl nitrate a t -70” C. leading to the displacement of lithium by nitrate. The structure has been confirmed by elemental analysis and conversion to As an illustrative aminoferrocene. point concerning the quasiaromatic character of ferrocene, the nitroferrocene once prepared shows an infrared absorption a t 1507 cm.-l typical of aromatic nitrogroups. Ferrocene can be oxidized to the ferrocinium ion electrochemically in water (db), in 90% ethanol ( I S ) , and in acetonitrile (10). The electrochemical process has been shon-n to be reversible Present address, Electrochemical Research & Development Department, Bell Telephone Laboratories, Inc., Murray Hill, K. J. VOL. 35, NO. 12, NOVEMBER 1963
1871
in these solvents. A number of ferrocene derivatives have been electrochemically oxidized at platinum electrodes (7, 9, 10, d0-22) and/or reduced at mercury (9, 21, 22). The oxidations always proceed with the conversion of iron from the +2 to f 3 oxidation state and the potentials of the oxidations follow a Hammett substituent correlation. The reductions a t mercury are either the reduction of the ferrocinium ion to ferrocene or a reduction of the substituent on the ferrocene nucleus. I n general, the reductions of the substituents are similar in behavior to the analogous aromatic compounds, but the potential of reduction of a substituted ferrocene is always more negative than for the corresponding substituted benzene. The electrochemistry of nitroferrocene, which has not been reported previously, is of interest in several respects. First, aromatic nitrocompounds have been shown to yield a variety of electrochemical products depending upon reactivity and solution conditions (1). It is useful to investigate electrochemical methods of preparing the partially reduced species: nitroso-, hydroxylamino-, and hydrazoferrocene. Second, the unusual methods required for preparation of the compound suggest that the nitro group may mark the "borderline" of reactive differentiation between the aromatic compounds and the ferrocenes. EXPERIMENTAL
All polarograms were recorded with a Sargent Model XV Polarograph modified to have a onesecond pen response. The dropping mercury electrode was constructed of Corning Marine Barometer tubing connected to the usual standtubeleveling bulb arrangement. Typical capillary characteristics a t -0.50 volt us. SCE in 25% v./v. ETOH-H20 a t 60 cm. height of mercury were: m = 1.765 mg. per second, and t = 4.86 seconds. For all the data reported in Apparatus.
Table 1. PH 4.05 4.99 5.28 6.47 7.03 7.09 7.28
this article, the currents were measured as the maximum excursion of the recorder pen; capillary characteristics were recorded a t the potential of measurement in the solution in question. Controlled potential coulometric analysis experiments were performed in a cell similar to that described by Lmgane (11). The potentiostat employed has been described previously ( 2 ) . Reagents. All electrolytes were reagent grade unless otherwise specified. Nitroferrocene and the sodium salt of 1-nitrocyclopentadiene were furnished by K. L. Rinehart of these laboratories. Nitrogen which had been successively passed through a solution of vanadous sulfate and a portion of the electrolyte solution was used to remove oxygen. All polarographic solutions were 0.080M in total buffer, 0.20,M in KC1 and 25% (volume) ethanol. In all cases, the ratio of buffer components was between 1 to 4 and 4 to 1. For experiments where measurements were made anodic of 0 volts us. SCE NaC104 or NaNOa was used instead of KCL. Because of the small quantities of nitroferrocene available, polarograms were obtained from minimum volumes of solution contained in a small H cell (about 10-ml. volume). Cell resistance (1000 ohms) was sufficiently high that potentials have been corrected for IR drop. The temperature of all test solutions was 25.0' rt 0.5" C. RESULTS
Polarography of Nitroferrocene. Solutions of nitroferrocene in well buffered 25% aqueous ethanol show one well defined, diffusion-controlled wave in the region +0.4 to -2.0 volts us. the saturated calomel electrode (SCE). No anodic wave is observed a t the D M E in the same region. The limiting current is proportional to the square root of corrected mercury height with zero intercept (8). If freshly prepared solutions of nitroferrocene are examined as a function of time, it is observed that the current of
Diffusion Current Constant and Half-Wave Potentials for Nitroferrocene in Buffered 25% v./v. Ethanol-Water Solutions E112 US. SCE, C X lo', iOb..,' iCW,.,b volt Molar (Pa.) fia. I 9.32 3.45 3.45 0.195 -0.455 5.89 9.23 0.335 5.51 -0.515 5.83 9.16 0.335 5.58 -0.530 5.58 9.07 -0.597 0.323 5.5s 9.93 15.2 15.2 0.809 Maximume 3.65 9.86 -0.625 0.195 3.34 9.13 6.96 Maximum. 8.97 0.518
a
Corrected for residual current.
c
Slight maxima were observed in the polarograms for the most concentrated solutions.
* Corrected for photochemical decomposition. 1872
0
ANALYTICAL CHEMISTRY
E VS S C E
Figure 1. Polarograms of nitroferrocene in 25% EtOH b y volume at pH 6.47 after successive exposures to light. Successive curves are displaced by 0.6 volt
the initial wave decreases rapidly while
a second, more cathodic process appears and grows with time. However, solutions prepared and examined under conditions of diminished light similar to that of a photographic darkroom show no change in diffusion current over a period of several hours. Figure 1 shows polarograms of a solution of nitroferrocene which was prepared in a darkened room and kept under such conditions with periodic exposure to a 100-watt incandescent bulb. With each exposure, the current of the second process increases a t the expense of the current of the first wave. I n the intervals between exposures to light, no change was observed. Figure 2 shows the change in current for each process and their sum as a function of cumulative exposure time under otherwise constant conditions. The change in each wave is a linear function of time. When identical solutions are maintained a t 0' C. and 25' C. in a lighted room, the rate of current change is the same for each solution. I n a phosphate buffer of pH 6.47, the ratio of the slope of current change with time for the first decreasing wave to that for the second increasing wave is -0.65, and a t pH 7.98 in trishydroxymethylaminomethane buffer the value is -0.67. The constancy of this value over the pH range examined allows a correction for the amount of decomposition which takes place in solutions exposed to light. Table I is a compilation of electrochemical characteristics obtained under several conditions of pH, concentration, and degree of decomposition. The pH range was 4.0 to 8.6, the concentration varied over a fourfold range, and some solutions decomposed to the extent of 20% during examination. A diffusion current constant, Z (8) from the simple Ilkovic equation
-1 I
1
FIRST A N D SECOND WAVES E-130V
VS S C E
between i and id, the plots are somewhat curved. If the curvature is removed by arbitrary selection of a n id, the lower portions become curved in much the same fashion one finds in firsborder kinetic plots. Although this behavior would suggest that the curvature is due to errors in measurement, the results of controlled potential electrolysis (vide infra) indicate that the process in fact may be more complex than a six-electron quantitative conversion to aminoferrocene. Below pH 6.47, the polarographic wave becomes distorted and assymmetric; the log plot for pH 4.05 is included in Figure 3 for comparison. The slopes for the straight line segments of each plot are: pH 4.05, 48 mv.; pH 6.47, 38 mv.; pH 7.09, 44 mv.; pH 8.16, 45 mv.; and pH 8.60, 43 mv. When compared to the expected value of 59 per n mv. for a reversible process, n is 1.3 and the process is definitely irreversible. Solutions of nitroferrocene in various buffers-e.g., phosphate, acetate, and t ris hydroxymethylaminomethane-produce similar behavior independent of buffer components. A plot of halfwave potential us. pH over the interval of pH 4 to 9 is linear with a slope of 58 mv. per pH. The extrapolated value of the half-wave potential a t pH 0 is -0.220 volt us. the SCE. Preliminary controlled potential coulometry experiments in trishydroxymethylaminomethane buffers indicate an TI value of the order of 5 when corrections are made for the small quantities of unreduced or photochemically decomposed nitroferrocene. Polarography of 1-Nitrocyclopentadienide Ion. I n the previous section, i t was noted t h a t nitroferrocene in solutions exposed to light decomDoses to give - a second electroactive species. Evidently, nitroferrocene is very sensitive t o light because rigorous efforts to purify the compound in the complete absence of light always resulted in a small (less than 1%) cathodic “decomposition wave” even when the melting point agreed with or was greater than the reported value. Since the decomposition did not seem to be sensitive to temperature or to pH, it was assumed that the reaction was not the result of chemical attack on the nitroferrocene. Consideration of simple photodecomposition schemes suggested that the product was nitrocyclopentadiene or its anion. Nitrocyclopentadiene and its anion can each be written in three isomeric forms, presumably in equilibrium, of which one (the 1-nitro form) also has an aci structure. On steric and conjugative grounds (stabilization by formation of this aci anion) the 1nitrocyclopentadienide terminology is preferred, although none of the experiments in this work give any evidence
oK-J T I M E IN SECONDS
Figure 2. Diffusion current of a p H 6.47 nitroferrocene solution as a function of time when exposed to light
was calculated using the corrected values for the current and the appropriate values of masf>flow rate of mercury and drop times. The relative standard deviation of these data was 4.7%. This uncertainty is satisfactory since stock solutions contained only 15 to 20 mg. of nitroferrocene and two dilutions were required prior to polarographic examination. No significant trend in the diffusion current constant was found as a consequence of changes in pH or concentration of depolarizer. When the overall number of electrons was evaluated using the diffusion current data and reasonable choices for model compound diffusion coefficients such as nitrobenzene or nitr ,naphthalene, the I1 9 , n values were inconclusively nonintegral. A more direct and reliable estimate was obtained from comparison of the diffusion currents for the reduction of nitroferrocene and the oxidation of the parent ferrocene which undergoes a reversible one-elec tron oxidation to ferrocinium ion (IO, IS, 22). The results of the comparison are shown in Table 11. Because the ferrocene solubility is limited in ttqueous solutions, the alcohol content for both compounds was increased to 50% by volume. The ratio of diffusion curr ?nt constants thus obtained was 6.02, with a relative Since standard deviation 3f 3.5% ferrocene in these solvents produces a one-electron oxidaticn, the ratio requires a six-electron reduction for the nitroferrocene. On this basis, the expected product is the aminoferrocene. Plots of log [(id -i)/i] us. potential for the first polarographic wave as a function of pH are shown in Figure 3. At p H 6.47 and greater, the curves are essentially linear over a ratio of i/id of 1000. At the more negative potentials corresponding to veq. small differences
I -0 5
-2
-0 3
Figure 3.
I
I
-0 7 E V S S C E
I
-0 9
- i)1 vs. L i l
Log r ( i d
E for
buffered 25% EtOH b y volume solutions of nitroferrocene A.
B. C. D. E.
pH pH pH pH pH
4.05 6.47 7.09
8.19 8.60
for this particular configuration of the electroactive species. A sample of 1nitrocyclopentadiene sodium salt was obtained and examined polarographically under conditions similar to those for the nitroferrocene. I n buffers more basic than pH 7.3, 1-nitrocyclopentadiene shows one well defined diffusion-controlled polarographic wave. The half-wave potential of this compound is identical within experimental error with the second decomposition wave for nitroferrocene.
Table II. Diffusion Current Constants for Ferrocene and Nitroferrocene in 50% v./v. EtOH/H20 Solutions
Concentration,
I
mMa
Ferrocene 0.488 0.343 0.316 0.418 0.252 0 . 467b
1.23 1.02 1.19 1.06 1.29 1.22 7 = 1.16“
Nitroferrocene 0.558 0.459 0.418
I
7.16 6.72 7.08 = 6.99
@Each solution waa prepared by independent weighing. * 0.1M HNOI used as supporting electrolyte rather than 0.1M NaNOs. c A value of 1.27 is reported by Page and Wilkenson ( l a )in 90% EtOII.
VOL 35, NO. 12, NOVEMBER 1963
. I
1873
The half-wave potential extrapolated to p H 7.00 is -1.14 volts us. SCE and shifts 90 mv. per p H unit for both nitrocyclopentadienide and for the decomposition wave. The icatter in the results for the decomposition wave is greater than that for the nitrocyclopentadienitlP because the extent of deconipozitiun was usually not large; diffusion currents were normally less than 0.5 pa. Electron Spin Resonance Spectrometry. The use of e.s.r. spectrometry t o identify free radicals produced electrochemically was first used by Geske and Maki in investigations of in situ generated nitrobenzene anion radicals in acetonitrile ( l a ) . Subsequent investigators have reported the spectra of similar radicals produced in aqueous solution (16, 16). Such studies are of particular interest in the present case since evaluation of coupling constants and spin densities yields useful information concerning the electronic character of the free radical. Solutions of nitroferrocene in 0.16N trishydroxymethylaminomethane or in unbuffered KCl in water produced an e m . spectrum upon prolonged cathodic electrolysis a t currents between 100 and 500 pa. The electrode potential was controlled by using a three-electrode potentiostat (5). During electrolysis the potential of the working electrode was maintained slightly anodic of the potential of the polarographic half-wave. Prolonged electrolysis, however, did not improve the spectrum beyond a signalto-noise ratio of 4 to 1. The spectrum had three lines of line width 3 gauss separated by 15.1 gauss with an intensity ratio of 1 to 1 to 1. No additional hyperfine splittings were observed. This equiintensity triplet is assignable to nitrogen. Though not precisely evaluated, the “g” value of this radical is not significantly different than that for the free-electron. The radical observed during the electrolysis of nitroferrocene is believed to be the anion radical analogous to the species produced during the electrolysis of metachloronitrobenzene ( 5 ) . In other studies of radicals produced from aromatic nitro compounds in aqueous solution (1, 15, 16), the nitrogen splitting constants fell in the range of 11 to 13gauss. Piette, Ludwig, and Adams (15, 16) reported splittings of 23 to 25 gauss for the radicals produced by electrolysis of aliphatic nitrocompounds. Geske and Ragle (S) observed a n increase in the nitrogen coupling constant as the coplanarity of the nitro group with the ring was distorted. This distortion or steric hindrance ~5 as caused by increased substitution on the benzene ring. Out-of-plane distortion effectively decouples the nitro group from the ring, causing greater localization of the electron on the nitro group. Thus, radicals
1874
ANALYTICAL CHEMISTRY
resulting from unhindered nitro groups attached to aromatic nuclei should have coupling constants in the range of 11 to 13 gauss while radicals with a nitro group which is highly hindered or attached to aliphatic nuclei should have coupling constants approaching 23 to 25 gauss. The nitroferrorene radical is, therefore, similar to unhindered benzeneoid substances, although the slightly larger coupling constant observed implies a slightly less aromatic character than is present in the nitrobenzene anion radical.
Comparison of the polarography of nitroferrocene with other nitrocompounds reveals differences in several respects. First, nitroferrocene shows only a single wave over a p H range in which the aromatic nitrocompounds exhibit a change from two waves in acidic solution. These are ascribed to successive reduction of ArNOz to ArNHOH and ArNHOH to ArKH2 to the single wave in alkaline buffers ascribed to the total reduction of ArKOz to ArNH2. Second, the wave for the nitroferrocene, although irreversible, is much less so than for nitrobenzenes. plots The slope of the log [(id-i)/i] ranges from 38 to 48 mv. per decade for nitroferrocene compared with an average value of 70 for substituted nitrobenzenes ( I ) . When experiments are carried out in diminished light, current-time curves for a growing drop indicate that adsorption is not a n important process. Polarographic maxima are not observed except in concentrated solutions or in solutions where significant photochemical decomposition has taken place. Aromatic nitrocompounds in the same concentration range and in similar solvents show pronounced maxima (8). Finally, the overall reduction over the p H range studied must be a six-electron reduction to the amine. This conclusion is based on the comparison of the cathodic reduction currents for nitroferrocene with the anodic ferrocene-ferrocinium ion wave. Since the process is clearly irreversible, the pH-potential value of 58 mv. per pH, which is temptingly close to the theoretical value for a reversible process that consumes the same number of protons as electrons, can only be used to support the generality that hydrogen ion is required in the process at some point and that the product is probably electrically neutral. On the above observations, the overall process can be described by
+ 6e + 6HC (Cp),Fe-NHZ
ACKNOWLEDGMENT
The authors thank C. S. Johnson, Jr., for his valuable assistance in the e.s.r. phase of this work.
DISCUSSION
(Cp)?Fe-N02
and the low values of n from macroelectrolysis experiments indicate that the electrochemical mechanism cannot lie as straightforward as shown in Equation 2. Since it i> possible to observe a free radical, a t least one electron tranzfer step must occur prior to the don ebt electron tranbfer step. A general mechanism for aromatic nitrocompound reduction based on similar evidence has already been proposed ( 1 ) .
-+
2H@ ( 2 )
The nature and existence of the free radical observed in e m . experiments
LITERATURE CITED
( 1 ) Clark, R. P., Ph. D. thesis, University of Illinois, Urbana, Ill., 1961.
(2) Enke, C. G., Ph.D. dissertation, University of Illinois, Urbana, Ill., 1959. (3) Geske, D. H., Ragle, J. L., J . Am. Chem. SOC.83,3532 (1961). ( 4 ) Grubert, H., Rinehart, K. L., Tetrahedron Letters 1959, 1617. (5) Hartley, A. M., Visco, R. E., J. Electrochem. Soc., in press, paper 175 New York Meeting of The Electrochemical Society, 1963.( 6 ) Helling, J. F., Shechter, H., Chem. and Ind. 1959, 1157. (7) Hoh, G., McEwen, W., Kleinberg, J., J . Am. Chem. SOC.83,3949 (1961). ( 8 ) Kolthoff, I. ,,M, Lingane, J. J., “Polarography, 2nd ed., pp. 63 and 85, Interscience, New York, 1952. (9) Komenda, J., Tirouflet, J., Compt. Rend. 254,3093 (1962). (10) Kuwame, T. K., Bublitz, D. E., Hoh. G.. J . Am. Chem. SOC.8 2 , 5811 (1960). (11) Lingane, J. J., “Electroanalytical Chemistry,” 2nd ed., p. 461, Interscience, New York, 1958. (12) Maki, A. H., Geske, D. H., J . Chem. Phys. 30,1356 (1959). (13) Page, J., Wilkenson, G., J . Am. Chem. Soc. 74,6149 (1952). (14) Pauson, P., Quart. Rec. London 9, 391 (1955). (15) Piette, L. H., Ludwig, P., Adams, R. N., ANAL.CHEY.34,916 (1962). (16) Piette, L. H., Ludwig, P., Adams, R. N.. J . 4 n t . Chem. SOC.83, 3909 (1961): (17) Rausch, M. D., Fischer, E. O., Grubert, H., Ibicl., 82, 76 (1960). (18) Rausch, XI. D., S’ogel, M., Rosenberg, N., J . Chem. Educ. 34, 268 (1957). (19) Schlogl, K., Osferr. Chem. Ztg. 59, 93 (1958). (20) Tirouflet. J., Boichard, J., Bull. SOC. ‘ Chim. France 1960, 4. (21) Tirouflet, J., Boichard, J., Ibid., p. 1032. (22) Tirouflet, J., Boichard, J., Conapt. Rend. 250, 1861 (1960). RECEIVEDfor review June 12, 1963. Accepted August 15, 1963. Division of Analytical Chemistry, 140th Meeting, .kCS, Chicago, Ill., September 1961. Taken in part from a dissertation presented by Robert Yisco t o the Graduate Faculty of the Cniversity of Illinois in partial fulfillment of the requirements for the degree of Doctor of Philosophy. Work was supported in part by the n’ational Institutes of Health, Grant Number RG649;. R. E. S’isco held an Electrochemical Society Summer Fellowship for much of the period in which this work was done.
