these expectations are realized. The polarograms of Figures 1 to 3 indicate that the peak potentials fqr the propyl amines vary in this manner. This holds for all aliphatic compounds examined. On the other hand, the peak for dimethylaminoacetonitrile, a tertiary amine, occurs at a more positive potential than that of any secondary amine examined. This reflects the strong electron withdrhwing ability of the CHICIV group. Correlation of polarographic half wave potentials and chronopotentiometric quarterwave potentials with the Hammett-Taft polar substituent constants (0and .*) has been demonstrated. However, rigorous agreement can only be expected for reversible electrode reactions. Correlations can be expected with irreversible reactions only when the transfer coefficients for all members of the group of compounds are the same. The pn, values in Table I1 indicate that most of the secondary and tertiary amines should be comparable, while the primaries should not be. This is reflected in the Rimilarity of slopes for the curves of dipropylamine and tripropylamine in Figures 1 to 3 as contrasted with the slope of the propylamine peak. These curves are typical of those obtained for members of each dass of aminei. Peak potentials are plotted in Figure 4 against the sum of u* ( 8 ) for corresponding compounds. h fair linear relationship is obtained for secondary and tertiary amines with a separate curve for the primaries. I t seems reasonable to speculate that these results, the similarity of pn, values for
BZ NH,
I
I
1.2
I
1.3 PEAK
Figure 4.
1.4 POTENTIAL
I
I
I
I
1.5
1.6
1.7
1.8
I
( V O L T S V S . N.H.E.)
Polar effect of substituents on peak potential
secondary and tertiary amines, and the linear n* relationship for secondary and tertiary, as compared with Primary amines. might be due to existence of two separate riaction mechanisms. In an attempt t80 throw some light on this question, a chemical examination of the react'ion Droducts from some amine oxidations- is being carried out in this laboratory. LITERATURE CITED
(1) Dapo, K. F., Llann, C. K., ANAL. CHEM.35, 677 (1963).
(2) Loveland, J. w.,Dimeler, G. R., 33p 'lg7 (1961). (3) Mann, C. K., Ihid., p. 1484. (4) Matsuda, H., Ayabe, Y., 2. Elektrochem. 59. 494 (1955). ( 5 j @Dinhefi,- J. F., c. K., \ - - - - /
ANAL. CHEM.
36$ 209i (1964).
(6) Russell, C. D., Ihid., 35, 1291 ( 1 963). ( 7 ) \ T ~Stackelberg, ~ M., SI., Toome. V.. 2. Elektrochem. 57. 347 (1953).' ' (8) Taft, R. W., "Steric Effects i n Organic Chemistry," SI. S. Newman, ed., chap. 13,.,Wiley, New York, 1956. RECEIVEDfor review May 1, 1964. Accepted September 22, 1964. Work supported by research grant GSI-10064 from the U.S. Public Health Service.
Voltammetric Behavior of the System NitrosobenzenePhenylhydroxylamine at the Graphite Electrode LAURA CHUANG, ILANA FRIED, and PHILIP J. ELVING The Universify of Michigan, Ann Arbor, Mich. The nitrosobenzene-phenylhydroxylamine couple behaves reversibly a t the stationary pyrolytic graphite electrode in 50% ethanolic buffered sohtion. The use of graphite allows the system to be studied a t pH as low as 1.6, compared to the limit of pH 4 with the dropping mercury electrode imposed by the oxidation of mercury a t the potentials involved. Potentials and currents at the graphite electrode are quite reproducible. The variation of the half-peak potential with pH is represented by E,iz = 0.33-0.060pH in agreement with previous polarographic and potentiometric measurements. The current varies linearly with concentration. Rotating the elec-
2426
s
ANALYTICAL CHEMISTRY
trode increases the current as expected but decreases slightly the agreement of the potentials for the reduction of nitrosobenzene and the oxidation of phenylhydroxylamine. Above pH 8, phenylhydroxylamine is rapidly oxidized chemically to nitrosobenzene. At pH 12, nitrosobenzene is also chemically altered.
- phenylhydroxylamine couple behaves reversibly both potentiometrically and polarographically (3, 12, 1 4 ) . Consequently, the voltammetric behavior of nitrosobenzene and phenylhydroxylamine a t the pyrolytic graphite electrode was
N
ITROSOBENZENE
investigated over a large range of p H t o obtain data over the pH range where the dropping mercury electrode cannot be used because of the oxidation of mercury a t the potentials involved; to compare the data obtained with those obtained for these two conipounds a t the dropping mercury electrode and potentiometrically; and to study the stability of the two compounds. The nature, theory, and applicability of the graphite indicating electrode has been recently reviewed (4). The stability and decomposition products of phenylhydroxylamine and nitrosobenzene were investigated because of their relevancy to the polarographic and voltammetric study and to
the analytical voltammetric determination of these compounds. Past work on the electrochemical behavior of the nitrosobenzene-phenylhydroxylamine system will be considered a t appropriate points in the subsequent discussion. EXPERIMENTAL
Reagents. Nitrosobenzene (Aldrich Chemicals) was recrystallized from 95% ethanol each time before use; its purity was satisfactory as judged by the agreement of its ultraviolet spectrum a n d melting point with t'he literature values. Phenylhydroxylamine was prepared, when needed, following Vogel ( 1 6 ) ; the crude product was recrystallized from ether. Stock solutions, 5 m N in phenylhydroxylamine or nitrobenzene, were prepared by dissolving the weighed amount of freshly prepared and recrystallized material in 957c ethanol. Double strength aqueous buffer solutions were prepared from reagent grade chemicals. The buffer systems used and their useful ranges of potential are listed in Table I. McIlvaine buffers could not be used above p H 5 because of precipitation of salts on alcohol addition. Nitrogen used for oxygen removal was passed through two ammonium vanadate solutions reduced by zinc in HC1, then through water, and finally through 5Oyc ethanol. Apparatus. Voltammograms were recorded on a Leeds & Northrup T y p e E Electro-Chemograph a t a polarization rate of 200 mv./minute; the water-jacketed H-cell (10) was maintained a t 25' & 0.1' C.; the saturated calomel electrode (S.C.E.) was separated from the electrolysis leg by a sintered disk and a n agar plug saturated with KC1. The pyrolytic graphite electrode (P.G.E.) is prepared by cutting a pyrolytic graphite block (kindly supplied by the General Electric Co.) into a cylinder, of 0.4-cm. diameter with the ab plane of the graphite parallel to the end used as the electrode surface. (The ab plane is that parallel to the layers of graphite; the ac and bc planes are, consequently, those perpendicular to the graphite layers.) The side of the cylinder is coated with a mixture of 6 drops of Epon Resin 815 (Shell Cheinicd Co.) and 1 to 2 drops of curing agent A (diethylaminopropylamine). The Epon resin mixture is also applied to the inside of one end of a 15- to 20cm. length of 6-mm. glass tubing, into which the P.G.E. is fitted. After thorough drying of the assembly in an oven at 110' C. for about an hour, the surface of the graphite inside the glass tubing is scratched with a pointed object until graphite powder appears; this serves for electrical contact. The end of the P.G.E. is polished with S o . 600 silicon carbide paper mounted on a rotating plate until a smooth, even surface is obtained. The electrode is held against t,he paper by a wooden guide designed. to keep the electrode
Table I.
Buffer Systems Used; pH and Decomposition Potentials a t Pyrolytic Graphite Electrode
Buffer composition KC1 HC1 McIlvaine McIlvaine McIlvaine NH4C1 NH, NHkC1 NH, NHAC1 NH, KCi KOH
+
+ + + +
Double strength 1.08 4.78
8.85 12.0
--
nHa r
Single strength 1.42 2.56 3.46 4.86
8.80
EtOH, 50% 1.63 3.34 4.22 5.42 6.60 7.90 8.58 12.53
Decomposition potentials,*volt Cathodic Anodic - 0 . 3 5 to - 0 . 4 2 0 86 '
- 0 76 - 1 08 to - 1 25 -0.58 -0.80 -1.01
0 81 0 70 t o 0 i l
0.52
pH values given refer to 1M ionic strength buffer prepared, its dilution by an equal volume of wat,er, and its dilution b y 95% ethanol as described for the preparation of the test solution. * Potential where the current in 50% aqueous ethanol solution deviates by 1 pa. from the straight-line portion of the background curve is taken as the supporting electrolyte decomposition potential. These potentials are us. S.C.E. and are reproducible to f 5 0 mv. a
perpendicular to the rotating silicon carbide Attempts to resurface the electrode on the lathe used to resurface waximpregnated spectroscopic graphite electrodes failed because of the breaking and splitting of the pyrolytic graphite. Coniparison of results from wax-impregnated electrodes resurfaced on the lathe and of pyrolytic graphite electrodes resurfaced by polishing showed that the surface area of the latter was as reproducible as that of the former electrodes. The glass tubing containing the electrode is filled with an appropriate amount of mercury and a piece of copper or nichrome wire is inserted. The electrode was rotated, when necessary, by a Caframo motor mounted on a sturdy iron frame. A Cenco Lab jack was used to raise and lower the H-cell to position the electrode. p H was measured with a Leeds I% Northrup Model 7664 p H meter, which was standardized with aqueous Beckman buffer standards of p H 4, 7, and 10. Because the p H values reported in this paper are those measured by a
Figure 1 . Variation of half-peak potentials of nitrosobenzene, phenylhydroxylamine, and related compounds in 50% aqueous ethanol
A Reduction waves observed in nitrosobenzene solution X Oxidation w a v e of phenylhydroxylamine 0 Reduction waves observed for basic solutions of phenylhydroxylamine 0 Azobenzene reduction 5 Azoxybenzene reductlon Roman numbers signlfy the w a v e order for a single pattern
conventional glass electrode-S.C.E. pair on 50y0 aqueous ethanol test solutions of ionic strength about 0.5111 (Table I), the exact interpretation of the p H is not clear, but the general pattern of the variation of behavior with p H is satisfactory. Spectrophotometric measurements were made with a Beckman Model D B recording spectrophotometer and stoppered 1-cm. quartz cells. Procedure. T h e voltammetric test solution was prepared immediately before use by pipetting 25 ml. of buffer solution into a 50-ml. volumetric flask, adding the appropriate amount of phenylhydroxylamine or nitrosobenzene stock solution, and diluting t o the mark with 95% ethanol. After mixing, the volume was again brought to the mark with 95% ethanol and remixed; a suitable volume was transferred to the H-cell. The mixture of equal volumes of water and 95Yc ethanol contracts about 2%. A standard procedure was followed in the use of the P.G.E. The P.G.E. was polished before each run, unless otherwise indicated, then rinsed with
?
PH
VOL. 3 6 , N O . 13, DECEMBER 1 9 6 4
2427
distilled water, dipped nioirientarily into 0.00570 Triton-X-100 solution, and rinsed with 50% ethanol and then with test solution. The electrode was inserted in the test solution, which had been purged with nitrogen for 10 minutes; passage of nitrogen was continued for 1 minute. The initial voltage was then applied for 0.5 minute, after which the potential scan was started.
oxidation of the phenylhydroxylamine to nitrosobenzene in the test solution, with the possible concomitant formation of azobenzene, and the slow oxidation of phenylhydroxylamine in its stock solution to nitrosobenzene and azoxybenzene. These reduction waves and their origin will be discussed in detail later. E,,z for nitrosobenzene reduction varies linearly with pH; E,!z for phenylhydroxylamine oxidation falls on the straight-line plot of E , us. pH for nitrosobenzene (Figure 1). The relationship between 13p,2 and pH for these two waves is represented by the following equation for t,he range of pH 1.6 t,o 12.5.
VOLTAMMETRIC BEHAVIOR OF NITROSOBENZENE A N D PHENYLHYDROXYLAMINE
The essential aspects of the voltammet'ric behavior of the reversible redox couple, nitrosobenzene-phenylhydroxylamine, a t the stationary pyrolytic graphite klectrode are summarized in Table I1 and Figure 1. Nitrosobenzene gives one reduction wave in the pH range of 1.6 to 12.5. A small second reduction ware (Eprp = -0.90) cauwd by azoxybenzene (2) appears a t pH 12.5, presumably as a result of t,he decomposition of nitrosobenzene (cf. subsequent discussion of stability). Phen.lhydroxylariiine gives one oxidation wave between p H 1.6 and 5.4. -it higher pH values, reduction waves may appear because of the rapid
d
~
E,.z
=
0.33
- 0.060 1)H
This result is in excellent agreement with previous measurements on these compounds at, the drop1)ing mercury elect,rode-e.g., Smith and \Taller ( 1 4 ) found for solutions-which were 10% ethanol and 0.00570 gelatin, between p H 4 and 10
El,* = 0.33 - 0.061 pH
20
(1)
(2)
Elofson and .itkinson (3) found for
60
sb
,A0
TIME, m l ~ .
,i
Figure 2. Variation of ultraviolet absorption of phenylhydroxylamine in 9570 ethanol with time in the presence and absence of air 0 In presence of air (0.02mM phenylhydroxylamine) X in absence of air (0.07mM phenylhydroxylomine) Numbers on the lines indicated the wavelength in rnp a t which the measurement was m a d e
reduction of nitrosobenzene in aqueous solution (pH 6 to 10) Table It.
PH 1.6 3.3 5.3 5.4 6.6 7.9 8.6 12.5
Voltammetric Data for Nitrosobenzene and Phenylhydroxylamine at Stationary Pyrolytic Graphite Electrode"
Phenylhydroxylamine Nitrosobenzene_ _ Wave I Wave IIb i,, pa. E,/*, volt i,, pa. E n l ~volt , i,, pa. -Ep!?,volt 2.60 0.240 -1.80 0.261 2.58 0.123 -1.91 0.123 1 73 -0.004 -1.56 -0.010 -1,i7 -0.081 2.88d -0 153 1.98 -0.187 2 . 0 S d -0.188 1 61 -0,412 0.94d -0.416 1.34 -0.785
Wave 1 1 1 7 ______ i,, pa. E,,z, volt
0.32 1.0s 0.80
-0.773 -0.757 -0.921
a Nitrosobenzene or phenylhydroxylanline concentration = 0.2mAI. For i,, positive sign indicates a cathodic or reduction wave, and minus sigti, an anodic or oxidation wave. * This wave is caused by azobenzcne produced from oxidation of phenylhydroxylamine as discussed in the text. c This wave is caused by azoxybenzene formed by oxidation of phenylhydroxylamine to nitrosobenzene, which reacts with tbe phenylhydroxylamine. d This wave is caused bv nitrosobenzene formed on rapid and complete oxidation of phenylhydroxylamine.
Table 111.
Concn , mdf
01 0 2 0 4 0 6 0 8 10 12 Mean Std. dev. 0
0 88 58
2 4 7 9 10
1s 20 20 24
0.339
0
(3)
and in 52% ethanol solution (pH 6 to 13) Eli2
=
0.323 - 0.0603 p H
(4)
These data are also in accord with the standard potential of 0.582 volt us. N.H.E. (0.337 volt us. S.C.E.) found for the system by potentiometric titration of each compound (12). Comparison of the E,,* values at the graphite electrode with the El values of Smith and Waller at the D.M.E. indicates that the graphite electrode values, on average, differ by about -9 mv. The slope of potential with pH at the P.G.E. is comparable with that obtained by Elofson and Aitkinson ( 3 ) in 52% ethanol solution; the slight
0 0 0 0 0 0
120 123 122 123 127 121
0 123 i.0 0024
8 12 10 12
8 9 4 0
1 1 $5 10 2
11 0 +14
72 64 44 00 1 1 25 1 4 7 10
At pH 12.5 E,,*, Volt
Phenylhydroxylamine at pH 3 . 3 EpD, volt "p/C
d C
ZP, Pa
412 425 429 423 429
8.6 11.6 12.4 12 5 11 3
-1.91
0.114
9.6
-5.74 -7.72 -9.26 -11.66
-0 424 f O 007
11.3 fl.0
0,119 0.113 0.lli 0.118 0 116 1 0 0026
9.6 9.6 9.3 9.7 9 6 1 0 15
-0 -0 -0 -0 -0
At 0.8and 1 .OmJl, humplike waves of 0 .5 pa, were observed at pH 3.3,extending from -0.45 to - 0.60.
2428
- 0.0584 p H
Dependency of Peak Current on Concentration of Nitrosobenzene and Phenylhydroxylamine
Sitrobenzene ___ At pH _______3 30 E , 2, volt Ip/C I , , Pa
zn,
=
ANALYTICAL CHEMISTRY
difference in potential at' pH 0 is actually a matter of how the line is drawn through the experimental points. .\ccordingly, one may safely conclude that the results obtained a t the pyrolytic gra1)hite electrode are comparable, if not identical. to those obtained with t'he dropping mercury elect,rode. I n addition, the pyrolytic graphit,e electrode allows evaluation of the nib-osobenzenephenylliydrox2-Iamiiie pH region below 4, where t'he D.M.E. cannot be used because of the oxidation of the mercury at the potentials involved. Smith and Waller (14) observed oxidation waves for phenylhydroxylamine at the 11.M.E. in the pH range of 4 to 10; under our experimental conditions, I)henylhydroxylamine is practically instantaneously oxidized to nit,rowhenzene at pH 7 or higher. JIixtureu of nitrosobenzene and pherivlliyti1.ox!-laiiiine were not run becaii.w voltaniiiio$rai,hing a compound which is a member of a reversible redox systcm. at a stationary electrode by apl)lying an initial potential a t which the compound gives an appreciable current, result's in a complicated current-potential pattern ( I ) , C o x c m m . ~ n o xDEPENDENCE.The dependencay of peak current, i,, and half-lmk potential, E , 2, on concentration a t pH 3.3 and 12.5 is summarized in. Tahle 111. The peak current varies essent'ially linearly with concentration for both compounds. In t'he case of phenylhydroxylamine, there may be some question concerning t8he purity of the pre1)arations used, as previously mentioned, and consequently regarding the absolute validit?. of the concentration and i, C values. The btandard deviation of .Eplz is about' 5 mv. in each case. EFFECTOF ELECTRODE ROTATION. To evaluat,e the possible increase in analytical sensitivity, the comparative voltaiiiinetric behavior of phenylhydroxylamine and nitrosobenzene was observed at stationary and rotating (123 r.p.m.) pyrolytic graphite electrodes. A l t1)H 3.3, 0.20111_11phenylhydroxylamine, which had i, = 2.43 pa. and E p 2 = 0.123 volt' a t the stationary electrode, gave the following values in three separate runs a t the rotating electrode: 5.50 pa., 0.143 volt; 5.72 pa.>0.136 volt; 5.62 pa., 0.148 volt. d t p H 1.6, 0.40mJI nitrosobenzene gar? values of 6.28 pa. and 0.245 volt at the stationary electrode, and 11.O pa. and 0.205 volt at the rotating electrode. E , values for the oxidation of phenylhydroxylamine and the reduction of nitrosobenzene at the rotated electrode are apparently less close to each other than those found a t the stationary electrode, because the devia-
tions are in the opposite direction to one another based on Figure 1. Because the current for both oxidation and reduction processes at a rotated graphite disk electrode has been shown (5, 6) to vary in the theoretically expected fashion with the square root of the angular velocity-i.e., rate of rotation in revolutions per minute, and directly with the electrode area-the current at other rates of rotation and electrode dimensions can be readily computed. STABILITY OF PHENYLHYDROXYLAMINE AND NITROSOBENZENE
Because phenylhydroxylamine is considered to be quite unstable, t h e extent of its decomposition under the polarographic conditions used was investigated to ascertain whether t h e changes would be significant. T h e results obtained, together with those from a less extensive study of nitrosobenzene, helped t o explain t h e reduction waves observed in alkaline test solutions of phenylhydroxylamine and the second reduction wave in pH 12.5 test solut.ion of nitrosobenzene. The complex series of chemical reactions possible between the various species in t,he sequence from nitrobenzene'to aniline as a result of electrochemical reduct,ion, which was originally
Table IV.
Time, minutes 0
27 52 88 106 916 0 15
35 50 80 120 0
22 43 63 91
investigated by Haber ( 7 ) , has been recently discussed by Harwood, Hurd, and Jordan (8). PHEXYLHYDROXYLAMINE. The experimental results, summarized in Table I V and Figure 2, are based on following the changes with time in the ultraviolet absorption spectra of solutions of freshly recrystallized and day-old samples of phenylhydroxylamine in ethanol, which were exposed to air (stock solut,ions were normally prepared in 95% ethanol) ; a similar solution of the same day-old sample in oxygen-free 95% ethanol, which solution was then kept sealed to exclude air; and a solution of the same fresh sample in air-exposed 957, ethanol, which was then converted into a typiciil test solution in p H 3.3 buffer . In 95% ethanol solution, phenylhydroxylamine shows absorption peaks at 236-8 and 280 nip. The absorbance a t 236-8 m p decreases with time, whereas that at 280 nip increases. With time, two more absorption peaks become apparent a t 305 and 220-226 nip; t'he former increases and the latter decreases with time. The latter is apparently masked during the initial period of time by the broad 236-8 nip band, upon which it first appears as a shoulder. Because nitrosobenzene in ethanolic media has two absorpt,ion peaks at
Variation in Ultraviolet Spectra of Phenylhydroxylamine under Different Conditions
Absorbance 280mp 236-8 mp 220-6 mp (A-1) Fresh sample (0.05mM) in 95yGEtOH 0 086 0 410 0 061 0 119 0 376 0 089 0 150 0 346 0 260 0 131 0 194 0 301 0 252 0 149 0 215 0 281 0 244 0 362 0 438 0 048 (A-2) Day-old sample (0.02mM) in 95YGethanol 0 048 0 330 0 070 0 307 0 037 n 090 n~ 286- _ 0 058 0 112 0 265 0 090 0 149 0 227 0 175 0 129 0 190 0 185 0 165
305 mp
218 mp
0 300
( B ) Day-old sample (0.07mM) in 95% ethanol in absence of oxygen 0 984 0 546
n
0 056 0 079 0 114
107
0 528 _-_
0 134 0 159 0 199
0 506 0 476 0 429
0 323
( C ) Fresh sample (0.069mM) in pH 3.3 McIlvaine buffera (305-315 mp) (275-285 mp) 0 0.034 0 092 0.438 30 0.038 0.096 0.432 59 0 046 0,098 0,430 99 0 051 0.102 0.420 117 0 051 0.102 0.420 900 0.146 0.146 0.360 a In pH 3.3 buffer, because of the strong absorption of the buffer system, the spectrum of phenylhydroxylamine cannot be seen below 235 mp; in addition, the bands at 305 and 280 mp are considerably broadened.
VOL. 36, NO. 13, DECEMBER 1964
2429
281.5 and 305.5 mp (15), it is apparent that phenylhydroxylamine in ethanol is gradually oxidized to nitrosobenzene. However, the rates of decrease of absorption at 236-8 n ~ pand of increase at 280 and 305 mp are not affected by the presence or absence of air-Le., air oxidation is not significant under the experimental conditions. Consequently, consideration was given to other possible oxidants for phenylhydroxylaniine Fvhich might be present. Phenylhydroxylamine is synthesized (16) in aqueous solution by the reaction C6H5XO2
+ 2Zn + H?O
-P
C6HsSHOH
+ 2ZnO
(5)
Stoichiometric amounts of nitrobenzene and Zn are used; no attempt is made to separate unreacted nitrobenzene. The degree of success of the react'ion depends on t,he efficiency of dispersing the two reactants-e.g., any Zn trapped by being coated by the ZnO formed would result in nitrobenzene being left. During t,he synthesis, droplets of nitrobenzene were apparent on the surface of the aqueous solvent, after t,he reaction of nitrobenzene and zinc dust was presumably complete; this nitrobenzene could react with phenylhydroxylamine to produce nitrosobenzene C6HsSHOH
+ CsHrh'Oz 2CsH&0
=
+ H20
(6)
I t is, therefore, possible that minute amounts of nitrobenzene, which were not removed from the phenylhydroxylamine during the recrystallization of the latter, were the oxidant. 1 week-old solution of phenylhydroxylamine in 95% ethanol showed three absorpt,ion peaks a t 324, 257-61, and 225-32 mp, which correspond to the peaks of azoxybenzene in ethanolic solut'ion, vie., 323, 261, and 231 mp (16). The formation of azoxybenzene by reaction of phenylhydroxylamine and nitrosobenzene is subsequently discussed. To minimize the oxidation of phenylhydroxylamine in its ethanolic stock solutions, the latt'er were prepared immediately before samples were withdrawn t,o prepare test solutions. In 50% ethanolic pH 3.3 buffered solution (Table IV) t,he decrease in phenylhydroxylamine after 106 minutes is only 4Cj, as compared with about 3270 in 9.5% et,hanol. This is most likely a pH effect because the rate of disappearance increases with increasing p H (cf. subsequent discussion). I t is also possible that in 95yc ethanol the alkalinity of the glass could aid the oxidation of the phenylhydroxylamine; the presence of an acidic buffer would prevent this. The voltammetric data (Table I) indicate that in test solutions of pH 7.9 or greater, phenylhydroxylamine is 2430
very rapidly oxidized, presumably by oxygen from the air, or dissolved in t'he test solution-e.g, within 30 minutes after solution peparation. h pH 7.9 buffered solution of phenylhydroxylamine had an ultraviolet absorption spectrum identical to that of nitrosobenzene in the same buffer system. This result is in agreement with the observation (11) that phenylhydroxylamine is rapidly converted in the presence of alkali to azoxybenzene; nitrosobenzene, as it forms, could react,with residual phenylhydroxylamine to form azoxybenzene ( I S , 16)
ANALYTICAL CHEMISTRY
CsHsNHOH f CsHsSO = CsHs-N=N-C6Hs
1
+ Ha0
(7)
0
The so-called third reduct,ion wave of basic solutions of phenylhydroxylamine, which are actually solutions of nitrosobenzene, is caused by the phenylhydroxylamine formed in the wave I reduction process reacting a t the electrode-solution interface with nitrosobenzene to produce azoxybenzene (Equation 7) ; some azoxybenzene might have been formed in the phenylhydroxylamine stock solution as previously mentioned. The azobenzene, whose reduction is the cause of the second reduction wave, might be produced by reaction of azoxybenzene and phenylhydroxylamine at high pH (azobenzene was not detected in solutions of pH lower than 12.5).
+
C ~ H ~ - - N J = ~ , T - C ~ H ~C6HJ'JHOH
.1
0 C6Hs--N=N-csHs CsHsNO
+ + HzO
+
(8)
The possible format,ion of azobenzene from the reaction of nitrosobenzene and aniline was not considered because the reduction product of nitrosobenzene under the present experimental conditions is phenylhydroxylamine and not aniline. Consequently, the presence of aniline in a nitrosobenzene solution or in a basic phenylhydroxylamine solution is not very likely. NITROSOBENZENE. The fact that the second nitrosobenzene wave only appears at times and then only at high pH rules out its origin in an impurity. At pH 12.5, nitrosobenzene absorbs at 218, 283. and 308 mp. The absorbance of a pH 12.5 buffered solution (5070 ethanol) of nitrosobenzene at 308 mp was followed with time; the absorptions at 218 and 283 overlap with t.hose of azoxybenzene at 231, 261, and 323 mg (15). In 90 minutes there had been a 13yc decrease in absorbance. Over a 4-hour period, the decrease in absorbance of nitrosobenzene in an equivolume ethanol-water mixture was only
2% and in a p H 1.6 buffered solution was nil. Consequently, the second reduction wave observed at pH 12.5 is caused by azoxybenzene formed by reaction of phenylhydroxylamine produced in the first x a v e process with unreduced nitrosobenzene, as previously d i m w e d . The above discussion regarding the origins of the reduction waves of phenylhydroxylamine solution and the extra reduction wave of nitrosobenzene solution a t high pH is easily confirmed. If the variations with pH of t'he reduction E P z values of azoxybenzene and azobenzene are plotted on the same graph with the E,'2 values observed for solutions of nitrosobenzene and phenylhydroxylamine, the second reduction wave observed in nitrosobenzene solutions and the t,hird reduction wave observed in basic solutions of phenylhydroxylamine coincide with the azoxybenzene plot,. The second reduction wave in basic phenylhydroxylamine solutions appears to be caused by azobenzene and the first wave, as already shown, by nitrosobenzene. DISCUSSION
The useful potential range of the pyrolytic graphite electrode in 5oy0 ethanolic buffered solutions is from about 0.9 to -0.4 volt at p H 1.6 and shifts with increasing pH in a positive direction to become 0.5 to -1.0 volt a t p H 12.5. These potential ranges emphasize the major advantage and the major limitation of graphite electrodes in respect to potential range as compared to the dropping mercury electrode. Redox processes, which occur at positive potentials compared to S.C.E., can be conveniently studied at the pyrolytic graphite electrode (P.G.E.); on the other hand, the low overpotential of hydrogen on graphite limits the negative potential range available. In the present study, the use of graphite electrodes permitted extension of the lower limit of the p H range studied froin 4, which was the lowest possible with the D.M.E.. to p H 1.6. The following conclusions can be drawn from the voltammetric study of nitrosobenzene and phenylhydroxylamine a t the P.G.E. The nitrosobenzene-phenylhydroxylamine couple is reversible a t the pyrolytic graphite electrode as well as at the dropping mercury electrode. Its E , 2-pH relationship in 50% ethanol, represented by the equation
Ep/Z = 0.33 - 0.060 p H
(1)
is in agreement with the DA1.E. polarographic (3, 14) and the platinum electrode potentiometric (12) data on the couple.
In 95% ethanol, phenylhydroxylamine is oxidized to nit,rosobenzene and can further react, with the nitrosobenzene thus formed to form azoxybenzene. .lzobenzene can be detected occa>ionally as an intermediate oxidat,ion product of phenylhydroxylamine cawed by react,ion of the latter with azoxyhenzene. The difference b e t m e n the reduction and oxidation E,] values of nit,rosobenzene arid phenylhydroxylamine, reh1)ectively. i:, larger at the rotating P.G.I
