give high results. By polarographic comparison (Figure 1 and Figure 3), iridium(1V) is as strong an oxidizing agent as is hypochlorite and both oxidants are capable of oxidizing palladium(I1) in the recommended supporting electrolyte, as verified by the result shown in Table 11. The potential selected is sufficiently positive that most of the possible interfering ions do not give waves, yet sufficiently negative that the excess hypochlorite is sensitively detected. Although palladium alone can be titrated with an error of less than 1%,the 3% criterion for interference was chosen because it is about the error level encountered in either the spectrophotometric method or the gravimetric method for the weight range selected. The conditions selected for the recommended procedure are about midway of the optimum chloride and azide concentration ranges. The chloride concentration used also aids in keeping lead in solution, an advantage if the palladium has been collected in a lead button by fire assay. The pH interval between 2.0 and 3.0 gives an optimum current response with respect to time and stability and is sufficient to prevent hydrolysis of many heavy metal ions.
The range between 0.5 and 2.1 mg of palladium was selected because the current stability is good, the variation in the ratio of hypochlorite topalladium is smal1,and large ratios of diverse ions to palladium are tolerated. Also, this range is slightly above that covered by spectrophotometry and below that conveniently amenable to gravimetry. To estimate the precision of the recommended procedure, eight titrations were made. An average ratio of 2.06:l was obtained, with a standard deviation of i 0.01 and a range from 2.05 to 2.08. Cerate, permanganate, and dichromate cannot be used as oxidants in place of hypochlorite. The first two reagents quantitatively oxidize azide to nitrogen in acid solution, and dichromate does not appear to oxidize palladium(I1) in the chloride-azide medium. RECEIVED for review November 16,1967. Accepted February 21, 1968. Work done under the auspices of the U. S. Atomic Energy Commission. Presented at the Eleventh Conference on Analytical Chemistry in Nuclear Technology, Gatlinburg, Tenn., Oct. 10-12, 1967.
Electrochemical Oxidation of Cupferron James G . Lawless and M. D. Hawley Department of Chemistry, Kansas State University, Manhattan, Kan. 66502 The electrochemical oxidation of cupferron is shown by electrochemical and spectroscopic methods to be a one-electron process which yields nitric oxide and nitrosobenzene as products. Evidence is presented for the regeneration of cupferron by a chemical reaction between nitrosobenzene and nitric oxide. Nitric oxide is also electrochemically active and can be oxidized stepwise to nitrate, the first step being the one-electron oxidation of nitric oxide to nitrite and the second step being the two-electron oxidation of nitrite to nitrate, In acidic solution, nitrosobenzene undergoes a two-electron reduction to phenylhydroxylamine, while in basic media coulometric reduction of nitrosobenzene is an overall one-electron process which leads to azobenzene. The latter product is shown to arise from a chemical reaction involving nitrosobenzene and its reduction product, phenylhydroxylamine.
tions of nitrosobenzene in basic solution, the electroc hemica oxidation of cupferron was investigated as a possible means for the in situ generation of this species. EXPERIMENTAL
THE AMMONIUM DERIVATIVE of N-nitroso-N-phenylhydroxylamine (cupferron) has been employed widely for the precipitation of metals from aqueous solution. In connection with end-point detection by electrical methods in these precipitation titrations, the electrochemical behavior of cupferron has been studied extensively at mercury electrodes ( I , 2). Although the polarographic results are not in total agreement, coulometric reduction of cupferron in acidic media has been shown to be an overall six-electron process which yields the corresponding hydrazine. In basic solution four electrons are involved with the final products being benzene and nitrogen (2). From an inspection of the structural formula of cupferron, we anticipated nitrosobenzene as a possible product of cupferron oxidation. Since we have been interested in the reac-
Instrumentation. A transistorized potentiostat-galvanostat of conventional design was used for most chronoamperometric and chronopotentiometric work (3). The detector for this unit was a Moseley Model 7030A x-y recorder. A second potentiostat was employed for experiments requiring larger voltage and current capabilities. The control amplifier of this circuit (ref. 4, Figure 15a) consisted of a Philbrick P45ALU operational amplifier in a noninverting configuration driving a Harrison Model 6824A amplifier. When this instrument was used as a coulometer, the readout voltage from the integrator was displayed on a strip-chart recorder. A manually operated potentiometer was used to control the potential difference between the working and reference electrodes. For rapid scan work, a Hewlett-Packard Model 3300 function generator equipped with a Model 3302 trigger served as the signal generator. The detector was a Tektronix Model 564 oscilloscope equipped with Type 2A6A and 2B67 plug-ins. Electrodes. A carbon paste electrode having a geometric area of 0.26 cm* was used as the working electrode for most chronoamperometric and all chronopotentiometric experiments. Details for the construction of this electrode and the renewal of its surface have been reported previously (5). A large platinum gauze served as the working electrode for controlled-potential coulometric studies at low pH. In strongly basic solutions, a mercury pool was employed as the working electrode for the reduction of nitrosobenzene. A small piece of platinum foil served as the auxiliary electrode for the chronoamperometric and chronopotentiometric
(1) I. M. Kolthoff and A. Liberti, J. Am. Chem. Soc., 70,l 885 (1948). (2) P. J. Elving and E. C. Olson, Zbid., 79, 2697 (1957).
(3) W. L. Underkofler and I. Shah, ANAL.CHEM., 35, 1778 (1963). (4) W. M. Schwarz and I. Shah, Zbid., p 1770. ( 5 ) C. Olson and R. N. Adams, Anal. Chim. Acta, 22, 582 (1960).
948
ANALYTICAL CHEMISTRY
0.4
E
- 0.4
Figure 2. Cyclic voltammogram of nitrosobenzene at pH 5 0.4
0 E
- 0.4
Figure 1. Cyclic voltammogram of cupferron at pH 5
1
Concentration = 2.02 X 10-3M; scan rate is 167 mV/sec; a McIlvaine buffer was used
+
0.1" c. Materials and Chemicals.
The organic chemicals which were studied electrochemically werc obtained from the following sources: cupferron, J. T. Baker, Baker Analyzed; nitrosobenzene, Aldrich Chemical Co. ; nitric oxide, Air Products; and azoxybenzene, Eastman, White Label. All of these were used as received. Because of the limited solubility of nitrosobenzene in aqueous solution, a small quantity of absolute ethanol was used to facilitate dissolution. All solutions of nitrosobenzene and cupferron contained 10% by volume ethanol. The pH values listed are apparent values. Reagent grade organic and inorganic chemicals were used in the preparation of buffer solutions. Prepurified nitrogen was used for deaeration of solutions. RESULTS AND DISCUSSION
Cupferron was observed to be unstable in solutions below pH 5 and to decompose to give a product whose oxidation wave interfered with the normal oxidation wave for cupferron. Consequently, the present study was limited to the pH range of 5 and above, where no appreciable decomposition occurred during the experiment. The cyclic voltammetric behavior of cupferron in the potential range of 0.7 to -0.5 V at pH 5 is shown in Figure 1 . On the first anodic sweep, a single oxidation wave is seen near 0.5 V. Although a cathodic wave corresponding to the reduction of the oxidation product back to cupferron is not observed after reversal of the potential scan, a reduction wave does appear near -0.3 V. Since cupferron itself does not give a reduction wave at this potential, this process must arise from a product of cupferron oxidation. The anodic wave that is observed at 0.3 V on the second and all subsequent anodic sweeps is seen only if the cathodic process at -0.3 V is first made to occur. In order to determine the chemical processes responsible for this behavior, the cupferron system was examined by current reversal chronopotentiometry. With a current program of i, = 10.828 i n / (i, and i, represent the magnitudes of the anodic and cathodic currents on the first cycle, respectively), the ratio of the cupferron oxidation time, tu, to the cathodic , found to be 1 or only slightly less than transition time, T ~ was
I
0
E
I
- 1.0
Figure 3. Cyclic voltammogram of cupferron in buffered solution Concentration = 3.75 X 1 0 - 3 ~ ; scan rate is 167 mV/sec; a pH 5.0 McIlvaine buffer was used
Concentration = 3.75 X lO+M; scan rate is 167 mV/sec; a McIlvaine buffer was used
experiments, while a mercury pool, separated from the working electrode compartment by a salt bridge, served as the auxiliary electrode for the coulometric work. The polarographic experiments employed a dropping mercury electrode of the conventional type with a drop time of about 5 sec. A saturated calomel electrode was used exclusively as the reference electrode. The cell was thermostated at 25.0
I
1.0
1 for values oft, varying from 1 to 15 sec (the ratio of rc/tcwas approximately 0.9 for tu equal to 10 sec). This result indicates that the ratio of the number of electrons in the anodic process to the number in the cathodic process is 0.5 (6). A comparison of the cyclic voltammograms for cupferron and a known sample of nitrosobenzene (Figures 1 and 2) suggests that the latter species is a product of cupferron oxidation. This has been confirmed by a comparison of an ultraviolet spectrum of an exhaustively electrolyzed solution of cupferron with the spectrum of the authentic compound. Coulometric and chronoamperometric reduction of nitrosobenzene at this pH gives phenylhydroxylamine (7), which accounts for the anodic wave that appears near 0.3 V on the second anodic scan. Since the reduction of nitrosobenzene is a two-electron process, the cupferron oxidation must be a one-electron process. +NNO -+ +NO I
+ NO + e
(1)
0-
+NO
+ 2H+ + 2e
+"OH
(2)
The number of electrons in the cupferron oxidation can be estimated also from chronoamperometric data. Assuming n to be 1, the diffusion coefficient for cupferron is calculated to cm2/sec, which is comparable to the value of be 8.2 X 8.81 x 10-6 cmz/sec reported previously from results for the polarographic reduction of cupferron ( I ) . Values of n equal to 2 or more would give results too low for the cupferron diffusion coefficient, If the direction of the potential scan is switched at 1.25 V rather than at 0.7 V, a second anodic wave appears on the first cycle near 1.2 V (Figure 3 ) . This oxidation process does not give rise to a corresponding reduction process on the cathodic sweep. In order to determine the numbers of electrons involved in the first and second anodic processes, chronoamperometric i t 1 1 2 values were measured at 0.64 and 1.2 V (Table I). Since the first anodic process has been shown above to be a one-electron oxidation, a comparison of i t 1 / *= 580 pA sec"2 at 1.20 V to ift-01/2= 141 KA sec1lzat 0.64 V suggests that the second anodic peak corresponds to a three-electron process. An electrochemical reaction consistent with this inter(6) T. Berzins and P. Delahay, J. Am. Chem. SOC.,75, 4205 (1953). (7) L. Chuang, I. Fried, and P. J. Elving, ANAL.CHEM., 36, 2426 (1964). VOL. 40, NO. 6, MAY 1968
949
4
OO
4
12
8
PH
E
Figure 4. Potential us. pH diagram for cupferron in buffered solutions
Figure 5. Cyclic voltammogram of cupferron at pH 9
For the pH range 2.2 to 8.0 McIlvaine buffers were used, and for the pH range of 9.2 to 11.0 Sorensen
Concentration = 3.72 X 10-3M; scan rate is 167 mV/sec; a sodium carbonate buffer was used
buffers were used
+ 2H20
+
NO3-
+ 4H+ + 3e
(8) E. Bamberger, Chem. Ber., 51,634 (1918). (9) A. N. Nesmeyanov and S. T. Ioffe, Zh. Obshch. Khim., 11, 392 (1941); Chem. Absrr., 35, 58692 (1941).
Table 1. Chronoamperometric Values of it1’* * for 3.42 X 10-3M Cupferron Time, sec 1.5 2.5 5.0 7.5
PH 5 E = E = 0.64 V 1.2 V 146 150 158 162
580 580 581 585
In PA secl/*,
950
ANALYTICAL CHEMISTRY
PH 8 0.64 V
1.2 V
0.1MNaOH E = 0.68 V
177 178 186 193
563 563 573 584
271 274 277 28 1
I
0 E
I
I
- 0.8
Concentration = 4.18 X 10T3M;scan rate is 167 mV/sec
(3)
The identity of the second anodic wave was established by a comparison of the cyclic voltammograms of authentic samples of NO and NOZ- with the cyclic voltammogram shown in Figure 3. At this pH, the oxidation of nitrite occurs at nearly the same potential as the oxidation of nitric oxide, and separate oxidation waves for these processes cannot be discerned. The chronoamperometric data presented in Table I indicate that the oxidation of cupferron is not diffusion controlled at an applied potential of 0.64 V. Whereas a constant value of it112 is expected for the process described in Equation 1, experimental values of were observed to increase from 146 pA sec1/2at 1.5 sec to 162 pA sec112at 7.5 sec. Although a portion of this increase can be attributed to a departure from conditions of semi-infinite linear diffusion, a chemical reaction which produces additional electroactive species subsequent to the initial charge transfer is also indicated. Controlled-potential electrolysis confirms a kinetic complication in the cupferron oxidation. Instead of obtaining n equal to 1, as predicted by the process described in Equation 1, nonintegral values varying from 1.3 to 1.5 were found. Several workers have reported a chemical reaction between nitric oxide and nitrosobenzene and have postulated the formation of a nitrosyl derivative of N-nitroso-N-phenylhydroxylamine (8, 9). This nitrosyl derivative should be immediately hydrolyzed in aqueous solution to yield nitrite and the original electroactive species.
I
Figure 6. Cyclic voltammogram of cupferron in 0.lMNaOH
pretation is the oxidation of NO, a product that is formed by the first anodic process (Equation 1). NO
I
OB
$NO
+ 2N0
+
%
+N-O(N0) N=O
$-N-N=O
I
0-
+ NO,- + 2Hf
(4)
Since the process which is described in Equation 4 does give rise to the original electroactive species, this reaction could account for the observed results. In support of this reaction sequence, we note that after the addition of nitric oxide to a solution of nitrosobenzene an oxidation wave appears at the same potential as the cupferron oxidation wave. If the process described in Equation 1 were both electrochemically and chemically reversible, then EDfor the cupferron oxidation wave should be independent of pH for all values of pH greater than its pK.. For values of pH below the pK, of cupferron, a linear plot of E us. pH having a slope of - 59 mV would also be predicted. Although the data obtained in the pH range 2 to 11 are consistent with such an interpretation, the absence of any cathodic wave corresponding to the reduction of the cupferron oxidation product back to cupferron indicates that the process described in Equation 1 is clearly irreversible (no cathodic wave could be observed at scan rates up to 100 V/sec). The fact that there is general agreement in behavior for this irreversible process and the completely reversible case could arise if the initial electron transfer were reversible and the rate of the subsequent chemical reaction fast and independent of pH. The value of 4.5 that is indicated in Figure 4 for the pK, of cupferron is in agreement with previously reported values of 4.11 (IO) and 4.28 (I). Upon scanning anodically at pH 9 (Figure 5), three separate peaks arise in the oxidation of cupferron. The anodic wave near 0.5 V again corresponds to the one-electron oxidation of cupferron to nitric oxide and nitrosobenzene. The two anodic waves at more positive potentials arise from the stepwise oxidation of NO to NO3-, the first step corresponding to the one-electron oxidation of nitric oxide to nitrite, and the second (10) P. J. Elving and E. C . Olson, J. Am. Chem. SOC.,78, 4206 (1956).
I
I
1.2
I
I
0.8
I
0.4
0
E
Figure 7. Cyclic voltammetric behavior of nitric oxide and nitrite In 0.1M NaOH; A , nitric oxide; B, nitrite
step resulting from the two-electron oxidation of nitrite to nitrate.
+ 20HNOz- + 20HNO
-
+
NOz-
+ HzO + e
(5)
NOa-
+ HzO + 2e
(6)
After reversal of the potential scan just prior to anodic background, a single cathodic wave is observed for the reduction of nitrosobenzene near -0.35 V. Reverse current chronopotentiometry shows that this reduction process is kinetically controlled. Whereas the ratio of the anodic electrolysis time (tu 5 T ~to) the cathodic transition time ( T ~for ) the reduction of nitrosobenzene was approximately 1 at pH 5, rc/tuwas observed to decrease slowly with increasing tu at pH 9. The observed decrease can be attributed to two different chemical reactions, the more important of which is the reaction between nitrosobenzene and its two-electron reduction product, phenylhydroxylamine (7). +NO
+ +"OH
+
+N=N+
J-
+ HzO
(7)
0.5 V (Figure 6). If nitric oxide is first oxidized to nitrite, as postulated above, then this wave should correspond to an overall two-electron process. Thus, the value of it'/* ( E = 1.2 V) at pH 5 where n = 4 should be twice as large as the corresponding value of ( E = 0.68 V) in 0.1M NaOH. The experimentally observed ratio of 1.9 is in good agreement with this prediction. The cyclic voltammograms of nitric oxide and nitrite are shown in Figure 7. The anodic wave near 0.5 V corresponds to the one-electron oxidation of NO to nitrite, while the more anodic wave corresponds t o the two-electron oxidation of nitrite to nitrate. Because this latter process occurs on the edge of background, a comparison of n values for the two anodic processes cannot be made directly by chronoamperometry. Upon sweeping cathodically, the main reduction wave at -0.4 V is followed by a smaller wave at -0.75 V. The first peak corresponds to the two-electron reduction of nitrosobenzene to phenylhydroxylamine, while the second peak arises from the reduction of azoxybenzene. Although the nitrosobenzene anion radical has also been observed in the electrochemical (21) and chemical (12) reduction of nitrosobenzene, a potential-step chronoamperornetric experiment shows that this radical arises from the coproportionation of nitrosobenzene and phenylhydroxylamine. In this experiment cupferron was first oxidized at an electrode potential of 0.7 V for 10 sec. The potential was then stepped to -0.60 V which was sufficiently cathodic to reduce the nitrosobenzene formed by the anodic process. By comparing the current for the reduction of nitrosobenzene at time t > 7 with the cupferron current at time 7 (7 is the switching time) (23), the cathodic process was observed to effectively change from n = 2 shortly after switching to n = 1 at long times. If the radical does arise from a coproportionation reaction, as the evidence above suggests, azoxybenzene could be formed by a mechanism involving the dimerization of the anion radicals (24).
+NO
0
Since azoxybenzene is reduced cathodically of nitrosobenzene the effect of this reaction is to remove electroactive species from solution, which causes a decrease in the cathodic transition time. Although the process described in Equation 7 is relatively slow at pH 9, the reaction goes nearly to completion during the time required for the coulometric reduction of nitrosobenzene. In contrast t o n = 2 at pH 5 for the reduction of nitrosobenzene to phenylhydroxylamine, coulometric reduction of nitrosobenzene at pH 9 gave n = 1 and azoxybenzene. The latter compound was identified by both cyclic voltammetry and ultraviolet spectroscopy. The second reaction involves the products of the first electrode process (Equation 4) and has been discussed above. This reaction probably occurs more slowly than the first reaction, as evidenced by the nearly 1 :1 ratio of the anodic electrolysis time to the cathodic transition time for the chronopotentiometric data at pH 5. The fact that the nitric oxide wave is only slightly anodic of the cupferron wave (Figure 5) precludes a quantitative measurement of the extent of these reactions by either chronopotentiometry or chronoamperometry. The nitric oxide oxidation wave continues to shift cathodically with increasing pH. In 0.1M NaOH the anodic processes for nitric oxide and cupferron occur concurrently near
+ +"OH
24NO' OH
0-
24NO'
I
+N-N+
I
-0
+ 2H+
1
H 1
~ i+N-N-+ .
1
-0
F?
+N=Nr$
+ OH-
4
0
(9) The kinetics of the process which leads to azoxybenzene are currently under investigation. These results will be the subject of another paper. RECEIVED for review November 30,1967. Accepted February 15, 1968. Division of Analytical Chemistry, Midwest Meeting, American Chemical Society, November 1967. Work supported by the Bureau of General Research and the Research Coordinating Council of Kansas State University.
(11) P. B. Ayscough, F. P. Sargent, and R. Wilson, J. Chem. SOC., Sect. B, 1966, p 903. (12) C. J. W. Gutch and W. A. Waters, Proc. Chem. SOC.,1964, p
230. (13) W. M. Smit and M. D. Wijnen, Rec. Trao. Chim. Pays-Bas, 79, 5 (1960). (14) G. A. Russell, E. J. Geels, F. J. Smentowski, K. H. Chang, J. Reymonds, and G. Kaupp, J. Am. Chem. Soc., 89,3821 (1967). VOL 40, NO. 6, MAY 1968
951
