ELECTROCHEIMICAL OXIDATION OF LEAD ACETATE IN ORGANIC SOLVENTS CHARLES
W .
LEWIS
A N D
P A T R I C I A
C .
EDGE
Glass Research Center, PPG Industries, P . 0. Box 11472, Pittsburgh, Pa. 15238 The electrolysis of lead acetate in a variety of organic solvents yields a n amorphous anodic deposit which contains quadrivalent lead, absorbs strongly in the visible, and is a n electronic conductor. When transparent tin oxide electrodes are employed, the rate of dissolution of the anodic film may be followed spectrophotometrically. I n two such studies (dilute solutions of hydroquinone and concentrated solutions of Carbitol, both in tetrahydrofurfuryl alcohol) the rate of dissolution remained constant until a critical thickness was reached. Beyond this point the rate decreased continuously.
THEwork described was undertaken as, part of a program concerned with the development of light-absorbing devices operating over a wide range of optical densities (absorbances). The basic concept involved the electrodeposition of nonmetallic films on transparent electrodes. Hard durable films of tin oxide, pyrolytically deposited on glass, constitute ideal electrodes. Their use in electrochemical studies has been described (Kuwana et al., 1964). The most interesting results were obtained with the brown anodic film deposited from solutions of lead acetate in various organic solvents. Experimental
Electrodes were prepared by cutting slides (2.54 x 7.62 cm.) from an annealed commercial grade of tin oxidecoated plate glass. Electrolyses were performed a t 12.6 volts and an electrode spacing of 0.64 cm. Current and voltage measurements were made with a Keithley 610B electrometer. Optical d,ensities were measured a t 435 mp in a Beckman Model B spectrophotometer. Absorption measurements were made at 30.0" & 0.5" C. Solvents were distilled through a 4-inch Vigreux column a t atmospheric pressure with retention of a middle fraction boiling over a range of 2'C. or less. Anhydrous lead acetate was prepared by recrystallizing reagent grade trihydrate from glacial acetic acid and drying to constant weight in a vacuum oven a t 60" to 65" C. The hemisolvate, unstable above 56", was thereby avoided (Davidson and Chappel, 1933; Tarbutton and Vosburgh, 1932). T h e selection of a solvent for the electrodeposition required some exploratory work on the solubility characteristics of lead acetate. One or more hydroxyl groups in the solvent molecule seemed to be essential. Thus, ethylene glycol, diethy lene glycol, triethylene glycol, and their methyl and ethyl monoethers are solvents, whereas the diethers are not. Glycerol and its mono- and diacetates are solvents, whereas the triacetate is not. Dimethyl sulfoxide, an excellent solvent for lead acetate, is the only notable exception. Slight solubility in dimethylformamide was also observed. A striking dependence of the course of the electrolytic process on the nature of the solvent became apparent a t an early stage of the investigation. From a saturated solution in butanol, a uniform anodic deposit appeared after a few minutes, whereas no visibly discernible process took place a t the cathode, even after 16 hours of continu-
ous electrolysis. Similar behavior was observed with 1-pentanol, cyclohexanol, allyl alcohol, and phenethyl alcohol. On the other hand, when ethylene glycol, propylene glycol, glycerol, or ethanolamine was used, large crystals of lead appeared a t the cathode. At most, the anode showed a very slight darkening which did not appear to increase with time. Most of the other solvents tested (di-, tri-, and tetraethylene glycol, the methyl, ethyl, and butyl ethers of mono-, di-, and triethylene glycol and some 1,3-, 1,4-, and 1,5-alkanediols, etc.) yielded deposits a t both electrodes. The relative amounts, however, showed wide variations. Thus, with tetrahydrofurfuryl alcohol, heavy anodic deposits were accompanied by barely discernible lead crystals a t the cathode. With glycerol diacetate, the preponderant deposition took place a t the cathode. One might be tempted to generalize on the basis of these observations. At first it seemed that the simple alcohols promoted anodic deposition and the glycols promoted cathodic deposition. However, both benzyl alcohol and pinacol (supercooled) permitted deposition a t both electrodes. I n all solutions examined, the rate of increase of optical density was directly proportional to the current density. Thus the rate of darkening could be readily varied. Other lead salts were briefly examined. The benzoate and caprylate behaved very much like the acetate. However, the chloride, nitrate, p-toluenesulfonate, salicylate, sulfanilate, and p-hydroxybenzoate deposited lead a t the cathode only, regardless of the solvent. In water, the chloride and the nitrate, like the acetate, yielded deposits a t both electrodes. Since the current flowing in the film formation process did not decrease with time, we may conclude that the film material is an electronic conductor (Mackenzie, 1964). X-ray diffraction of a thick coating yielded a broad halo typical of an amorphous material. With thinner coatings the same result was obtained, except that the diffraction pattern of tin oxide was superimposed. The thick coatings shriveled upon soaking in water and could thereby be removed from the electrode. X-ray diffraction of this product revealed the presence of o-PbO2 in addition to the amorphous material described above. Chemical analysis gave the following results for a sample obtained by electrolysis of a 10% solution in tetrahydrofurfuryl alcohol:
Total Pb = 82.85 Pb(1V) = 60.03'; VOL. 8 N O . 4 DECEMBER 1 9 6 9
399
inhibit deposition. Obviously, there is a competition between the deposition and fading rates. I n cases where inhibition was observed, the anode showed a slight darkening which quickly reached a steady value and faded when the voltage was removed. Two systems were selected for spectrophotometric studies of the dissolution kinetics: Carbitol diluted with various amounts of tetrahydrofurfuryl alcohol, and a series of solutions of hydroquinone in the same solvent. The film was prepared by electrolysis of a lOLC solution of lead acetate in tetrahydrofurfuryl alcohol to a final optical density of about 0.9. The anode was then removed from the electrolyte, rinsed successively with tetrahydrofurfuryl alcohol and a separate portion of the test solution, and then placed in the filled optical cell in the spectrophotometer. The optical density ( D ) U S . time ( t ) curves were all qualitatively alike, and are exemplified in Figure 1. Initially, the optical density decreased linearly until a t some point ( t l , D1) the rate ( - d D ) d t ) began to fall off and thereafter decreased continuously and vanished a t (t?, D ? ) . The final stage of the reaction is shown on the expanded ordinate scale. The residual optical density, D,, may be attributed to absorption due to all sources other than the anodic film. The value of ( D l - D2) remained reasonably constant from one experiment to the next. For the Carbitol solutions it had the average value 0.55 f 0.02; for the hydroquinone solutions, 0.54 i 0.04. I n both series, the initial fading rate, ( D o - D l ) / t l ,was directly proportional to the concentration, c, of reagent (Figures 2 and 3). Furthermore, the time interval ( t 2 - t , ) was inversely proportional to c (Figures 4 and 5). Numerical values of the kinetic parameters characterizing the fading process are listed in Table I. The rate of the reaction, as far as could be ascertained spectrophotometrically, was not altered by gentle stirring of the liquid during the experiment. Variations of D O did not affect the rates nor the value of ( D l - D2). Working a t other wavelengths did not alter the observed kinetics, although the values of optical density were modified in accordance with the absorption spectrum of
If the film is looked upon as a mixed oxide, only 94% of its weight can be accounted for. The remainder was not identified. Upon standing in contact with the electrolyte, the anodic films showed an interesting range of stability. When butanol was used as the solvent, there was no detectable loss in optical density after 3 weeks. I n tetrahydrofurfuryl alcohol, a coating whose optical density was 0.9 a t 435 mF faded uniformly and was completely dissolved after 8 hours. With Carbitol and methyl Carbitol, the fading was complete in 45 to 50 minutes. The fading time was the same whether the cell was left in the open or shorted condition. Furthermore, removal of the cathode from the system did not alter the fading time. Next it was established that fading took place in the pure solvent as rapidly as in the lead acetate solution. Finally, it was found that the choice of solvent used for the electrodeposition did not have any effect on the fading process. Thus a coating formed in tetrahydrofurfuryl alcohol and subsequently transferred to Carbitol faded in 45 minutes. A coating formed in Carbitol and transferred to tetrahydrofurfuryl alcohol faded in 8 hours. The rapid fading in Carbitol and its methyl analog was rather surprising, in so far as the corresponding monoethers of mono- and triethylene glycol did not possess this property. I t seemed reasonable to assume that some impurity might be responsible, even though gas chromatographic analysis could detect only trace quantities of water in redistilled Carbitol. T o test this hypothesis further, Carbitol was refluxed with reagent grade PbO, for 2 hours a t atmospheric pressure. Large quantities of water and divalent lead were formed, but the purified solvent, recovered by distillation, showed no loss in activity. Traces of water enhanced the rate and a slight retardation was observed when carefully dried materials were employed. I t may be concluded that Carbitol itself is the active material in the fading process, and that water has a mild catalytic effect. A number of additives were tested to see if the dissolution process could be accelerated. Of those that were tried, hydroquinone and several of its homologs were the most effective. The addition of 0.5% hydroquinone to a 10% solution of lead acetate in tetrahydrofurfuryl alcohol inhibited the anodic deposition process (at 12.6 volts). If higher currents were used (either by increasing the lead acetate concentration or raising the applied voltage), higher concentrations of hydroquinone were required to
Table I. Summary of Kinetic Data
(Do - Di)/tiC, L . Mole ,+fin.-' Carbitol series Hydroquinone series
0.0088 9.3
>
I-
w
Y
0.6
d
0.5
U,
D, = 0 5 5
z
t
,
: 15
min
0 05
E 0,4L
0 04 0 03
0.3
0.2
tiL LI1
O.'00
20
40
12 = 155 m i n . ,
60
1
1
1
80
1
1
100
1
120
1
1
140
1
.
T
.
160
TIME (Minutes) Figure 1. Fading of anodic film in 2.21M Carbitol 400
I & E C PRODUCT RESEARCH A N D DEVELOPMENT
+ 0.0011 Zt
1.1
C ( t 2 - tll, Min Mole L.-'
291.0 j, 13.0 0.145 =k 0.005
0.28
r
0
0.06
-
-
P
0.05
In 0)
2
3
0.04
E
x
0.16
W
k O.O?
a W
2z
0.02
m
J
4
t z -
0.01
1
C I
2
3
4
5
6
I
7
2
3
4
5
6
7
8
CARBITOL CONCENTRATION (Moles/Liter)
CARBITOL CONCENTRATION (Moles/Liter)
Figure 4. Reciprocal of fading completion time in Carbitol solutions
Figure 2. initial iading rate in Carbitol solutions
the film. Experiments performed in a different instrument (Beckman DU) were in good agreement with those obtained in the Model B. Discussion
For the two reagents studied, there is a striking similarity in the kinetics of dissolution of the anodic film. I n both cases there is a narrow optical density range above which the rate is const ant and below which it falls steadily with time. I n both cases, this transition occurs in the neighborhood of 0.54. The initial rate is directly proportional to the concentration. The average rate in going from an optical density of 0.54 to 0 is also directly proportional to the concentration. I n view of these similarities, it would appear that a similar mechanism is involved in both cases. I n the case of Carbitol, reduction of the quadrivalent lead does not appear to participate in the rate-determining process. Thus ethylene glycol rapidly reduces lead tetraacetate, whereas its monoethers are much less effective (Criegee, 1931). Severtheless, ethylene glycol is less effective as a fading agent for the anodic film (60 minutes as against 45 for Carbitol). The addition of lithium acetate or magnesium acetate had no noticeable effect on the fading rate. Since these materials would certainly depress the activity of any ionic species, it follows
071
O0.06
0.05
0.02
0.01
t/ I
0
I 1
Figure 5 . Reciprocal of fading completion time in hydroquinone solutions
010, 0
c
Z
0.07
0.06
2
0.05
004-
0
0001
0002
0.003 0004
0005
0.006
0007
0008
0009
H Y D R O Q U I N O N E CONCENTRATION ( M o l e s I L i t e r ) Figure 3. Initial fading rate in hydroquinone solutions VOL. 8 NO. 4 DECEMBER 1 9 6 9
40 1
that no ionic mechanism for the rate-determining step is indicated. The failure of stirring to alter the solution rate suggests that the process is not limited by any redox potentials existing in the immediate vicinity of the electrode. The kinetics for the initial dissolution process may be represented by the equation:
dD dt
-=
-hoc
where the rate constant, ho, depends only on the fading agent employed. I n the second stage the rate constant decreases with tkiickness and therefore with D - D2
dD
-= dt
-hc
where h is a function of (D- D2).I t follows from Equation 1 that ho is given by the first column of Table I. Integration of Equation 2 gives:
(3) Since D2 represents the contribution to the optical density not due to the film, the right-hand side of 3 is just the reciprocal of the rate constant integrated over the film density between the limits of 0 and the critical value of 0.54. This should depend only on the reactant and not its concentration. The constancy of c(t2 - t l ) follows a t once. The decrease in rate after the critical density is reached is subject to a t least two interpretations: 1. The film thickness is not uniform. Therefore the optical density decreases linearly until the regions that were initially thinnest have been stripped down to the substrate. Thereafter these areas do not contribute to the rate. As more and more of the substrate is exposed, the rate gradually falls to zero. 2. During the deposition, the initial coating is deposited on a crystalline tin oxide surface which influences
its structure and composition. After a suitable thickness is attained, the influence of the substrate can no longer be felt and the deposit has a fixed composition thereafter. During fading, the uniform part of the coating dissolves a t a constant rate. After the thickness has decreased to the critical value (at which D = D1),a progressively less reactive surface is exposed. The first mechanism seems unlikely for two reasons: I n the first place, the coated electrodes are shiny. Optical and electron microscopy failed to reveal any appreciable topographic irregularities. Secondly, the observed kinetics is inconsistent with the assumption made. If the reactivity of the film were the same throughout, it can be readily shown that a reciprocal relationship should exist between the linear rate, (Do - D i ) / t l , and the completion time (t2 - t l ) , which should be independent of the nature of the reactant. From the data of Table I we can see that the product (Do - D l ) / t l x ( t z - t l ) has the value 2.56 0.43 for Carbitol and 1.35 i 0.21 for hydroquinone. Thus the kinetics must change in the case of a t least one of these reactants when the optical density goes through the value D1. I t is most likely that the kinetics changes in both cases, but not to the same extent. The second mechanism is also consistent with the observation that inhibition of anodic deposition by hydroquinone is accompanied by a slight darkening of the anode which does not increase with time regardless of how long the voltage is applied.
*
Literature Cited
Criegee, R., Ber. 64B, 260 (1931). Davidson, A. W., Chappel, W., J . A m . Chem. SOC.55, 4624 (1933). Kuwana, T., Darlington, R. K., Leedy, D. W., Anal. Chem. 36, 2023 (1964). Mackenzie, J. D., J . A m . Ceram. SOC.47, 211 (1964). Tarbutton, G., Vosburgh, W. C., J . A m . Chem. SOC.54, 4537 (1932). RECEIVED for review July 8, 1968 ACCEPTED September 20, 1969
SSOLUTION OF PLUTONIUM IN DILUTE NITRIC ACID F .
J.
MINER,
J .
H .
N A I R N ' , AND
J .
W .
B E R R Y
Rocky Flats Division, The Dow Chemical Co., P . 0 . Box 888, Golden, Colo. 80401
THEreprocessing
of plutonium metal uza conventional aqueous chemical methods requires a conversion from the metal to a nitrate solution. Once in solution, the plutonium can be purified by ion exchange or solvent extraction, then precipitated as an oxalate, fluoride, or peroxide, and reduced back to the metal. In some methods used in the past, plutonium metal was converted to a nitrate solution by dissolving it in concentrated nitric Present address, Chemistry Department, Rice University, Houston, Tex. 77006 402
I & E C PRODUCT RESEARCH A N D DEVELOPMENT
acid containing trace quantities of fluoride. However, this dissolution was slow. I n addition, explosions sometimes occurred, caused perhaps by reaction products, which could include hydrogen. A method used more recently for converting the metal to a nitrate solution requires burning the metal to the oxide and then dissolving the oxide in a concentrated nitric-hydrofluoric acid solution (Molen, 1967; Stevenson and Paige, 1968). This method requires an extra step and, like the dissolution of the metal, is slow.
