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GADDERAMACHANDRA RAOAND LOUISMEITES
The Voltammetric Characteristics and Mechanism of Electrooxidation of Hydroxylamine
by Gadde Ramachandra Raol and Louis Meites Department of Chemistry, Polytechnic Inatitute of Brooklyn, Brooklyn, New York
(Receizled June 20, 1966)
The voltammetric behavior of hydroxylamine a t both mercury and platinum electrodes has been investigated over a wide range of pH values. At mercury electrodes in strongly alkaline solutions (pH 2 13), the rate-determining electron-transfer step is a two-electron oxidation to nitroxyl, NOH, which can either decompose to give nitrous oxide or react with hydroxylamine to give nitrogen. As the pH decreases, a one-electron oxidation to HNOH begins to predominate; on dimerization and decomposition this yields nitrogen as the end product. At platinum electrodes, the behavior is complicated by the slow chemical reduction of the oxide film by hydroxylamine and by the differing rates of oxidation on massive, finely divided, and oxide-coated surfaces.
Introduction I t has been known for a long time that the electrolytic oxidation of hydroxylamine must occur by a very complex mechanism. Early studies summarized by Glasstone and Hickling2 showed that nitrous oxide, nitrogen, and nitrite are usually the principal products; that the proportions in which they are formed depend on the nature and potential of the anode and on the composition of the solution; and that traces of oxygen and nitrate may also be formed under certain conditions. Lingane and Jones,3 investigating the behavior of hydroxylamine as an anodic depolarizer for platinum electrodes, obtained an n value of 2.7 in the controlled-potential electrooxidation of an unbuffered solution of hydroxylammmonium chloride containing sodium chloride as the supporting electrolyte. Davis4 obtained n values varying from 2 to 4.6 by controlled-potential coulometry a t platinum electrodes a t various pH values between 1 and 13, whereas chronopotentiometry under similar conditions seemed to suggest n = 6; he associOn the ated the discrepancy with the Oxide trod€! surface. The work of Lord and Rogers‘ shows that similar complications arise at gold and graphite electrodes. studies Of the direct and There have been catalytic reductions of hydroxylamine at dropping mercury electrodes under various conditions, but only Vivarelli6 appears to have studied the anodic wave obThe Journal of Physical Chemistry
tained at pH values above 7. From its diffusion current constant and a thermodynamic interpretation of the pH dependence of its half-wave potential, he concluded that it represents a two-electron oxidation to hyponitrite 2NHzOH
+ 60H-
=
N20P
+ 6H20 + 4e
This conclusion is unjustified because the oxidation is not reversible, and it provides no clue to the understanding of the complex behavior described below.
Experimental Section Voltammograms were recorded with a locally constructed polarograph incorporating a strip-chart recording potentiometer having an undamped full-scale response time of 1 sec. Half-wave potentials and values of (1 - a)%,were obtained from maximum currents on polarograms recorded without damping; average diffusion currents were obtained by suitable condenser damping. Initial and span potentials were (1) This paper is based on a thesis submitted by G. R. Rao t o the faculty of the Polytechnic Institute of Brooklyn in partial fulfillment of the requirements for the degree of M.S. in Chemistry, June 1966. (2) S. Glasstone and A. Hickling, “Electrolytic Oxidation and Reduction,” Chapman and Hall, Ltd., London, 1935. (3) J. J. Lingane and S. L. Jones, Anal. Chem., 23, 1804 (1951). (4) D. G . Davis, ibid., 35, 764 (1963). (5) S. S. Lord, Jr. and L. B. Rogers, ibid., 26, 284 (1954). (6) 8 . Vivarelli, Ann. Chin. ~ p p i .41, , 415 (1951).
ELECTROOXIDATION OF HYDROXYLAMINE
measured with a Rubicon precision potentiometer. Polarograms were generally recorded with a span potential of only 0.5 v, and the half-wave potentials obtained from them are believed to be reliable to within 2 mv. The platinum wire electrode (Catalog No. S30420) and 600-rpm synchronous rotator were obtained from E. H. Sargent & Co. The potentiostat and current integrator were obtained from Analytical Instruments, Inc. (Wolcott, Conn.), All accessory equipment and all techniques of manipulation and measurement conformed to good current practice.' Voltam0.02'; metric measurements were made a t 25.00 controlled-potential electrolyses were performed at 23 f 2". An approximately 0.1 F stock solution of hydroxylammonium sulfate was prepared and standardized by the procedure described by Vogel18 which involves oxidation with excess bromate in 1 F hydrochloric acid and iodometric determination of the excess. The titer of the stock solution decreased approximately 0.15% per month. Nore dilute solutions were prepared by volumetric dilution as they were needed, and mere never stored for more than 2 days. Analytical grade reagents and distilled water were used throughout. The sodium hydroxide solutions were essentially carbonate-free, though small concentrations of carbonate were found to be without sensible effect. Although hydroxylamine is known to decompose spontaneously in strongly alkaline solutions, measurements of the rate of decrease of its wave height showed that the extent of this decomposition mas inappreciable over the duration of any of the experiments described here.
Results and Discussion Oxidation at Mercury Electrodes. The half-wave potential of the anodic wave of hydroxylamine becomes more negative as the pH increases. In 0.1-10 F sodium hydroxide the wave is quite well defined, but its plateau is not quite exactly parallel to the residual-current curve, so that the diffusion current measured at the half-wave potential by the common extrapolation technique is a few per cent smaller than that measured in the center of the plateau. All values reported below pertain to the midpoint of the plateau and are corrected for the separately measured residual current. The slope of the plateau becomes more pronounced as the pH decreases, and at the same time the plateau becomes shorter as the wave encroaches on the initial current rise. At p H 9 the wave is just amenable to reasonably precise measurement. I n a phosphate buffer of pH 6.5 only the foot of the wave can be recorded; after increasing to about one-third of the expected diffusion current, the current falls abruptly
3621
toward the residual-current line. Similar behavior has been observed for many other oxidizable species in phosphate b ~ f T e r s , ~ and ~ ' ~ "LinganeIob has attributed it to the formation of a monomolecular film of mercury (I) phosphate. The height of this wave is diffusion-controlled. With 0.246 m F hydroxylamine in 0.1 F sodium hydroxide, for example, the value of id/hCOr1i2 was constant and equal to -3.21 f 0.03 pa cm-"'over a twofold range of values of h,,,, the height of the column of mercury above the capillary after correction for back-pressure effects. In the same supporting electrolyte, the diffusion current is proportional to the concentration of hydroxylamine between 0.09 and 2 mF. Over this range the mean value of i d / c was 11.52 f 0.16 pa mmole-1 1. for a capillary having m = 2.005 mg/sec and t = 4.12 sec a t the potential of measurement. A slightly lower value was obtained with 0.02 m F hydroxylamine, presumably because some air oxidation took place as the hydroxylamine solution was being added to the deaerated supporting electrolyte. The mean value of i d / c corresponds to a diffusion current constant of -5.72 pa mmole-' 1. mg-2'a sec"'. The diffusion current constant decreases as the concentration of sodium hydroxide rises: it is -4.71 in 1.0 F, -2.71 in 5.0 F, and - 1.39 in 10 F sodium hydroxide. Vivarelli6 reported I = - 2.82 in 4 F sodium hydroxide. The corresponding values of Iv"', 7 being the viscosity of the solution, behave in a strikingly anonialous fashion: from -5.78 in 0.1 F sodium hydroxide they decrease to -4.52 in 10 F sodium hydroxide. For hydrazine,ll an increase of nearly 50% mas observed over the same range of concentrations and was attributed to a decrease of the hydration number caused by the decreasing activity of water. Here, on the contrary, it appears that the changing activity of hydroxyl ion affects the average n value by altering the relative rates of the elementary processes involved. According to the IlkoviE equation, I = 607nD"'. Values of D,the diffusion coefficient of hydroxylamine, may be estimated from Jander's equation,12 taking the diffusion coefficient of ammonia as 2.36 X 10-5 cm2/sec in a dilute aqueous solution at 2 5 O . l ' Similar calcula(7) L. Meites, "Polarographic Techniques," 2nd ed, Interscience Publishers, Inc., New York, N. Y . , 1965. (8) A. I. Vogel, "A Text Book of Quantitative Inorganic Analysis," 2nd ed, Longmans, Green and Co., New York, N. P., 1951, p 375. (9) 0. H. hlUller and J. P. Baumberger, Trans. Am. Electrochem. Soe., 71, 169, 181 (1937). (IO) (a) G. D. Christian and W-.C. Purdy, J . Electroanal. Chem., 3, 363 (1962); (b) J. J. Lingane, i b i d . , 12, 173 (1966). (11) S. Karp and L. Meites, J . Am. Chem. Sac., 84, 906 (1962). (12) G. Jander and H. Spandau, 2. P h y s i k . Chem., A185, 325 (1939).
V o l u m e 7'0, N u m b e r 11
.Vovember 1966
GADDERAMACHANDRA RAOAND LOUISMEITES
3622
tions for hydrazine in a dilute solution of sodium hydroxide gave" n = -3.94, in agreement with controlled-potential coulometric data, which gave n = -4. For hydroxylamine, they give the results shown in Table I, which also includes the values obtained by
Table I : Polarographic and Controlled-Potential Coulometric n Values for the Oxidation of Hydroxylamine"
Supporting electrolyte
0.02 F borax1 F KNOI 0.01 F NaOH1 F KNOa 0 . 1 F NaOH 1 . O F NaOH 3 . 0 F NaOH 5 . 0 F NaOH 10 F NaOH
Polarographic
Coulometric
...
-1.354 =I= 0.025
...
-2.15
-2.49 -2.31
...
-2.15 -2.305 f 0.013 -2.020* 0.016
-2.27 -1.99
-1.77
...
The polarographic n values were computed from the diffusion current constants, using diffusion coefficients estimated as described in the text. The coulometric values were obtained in oxidations of 0.1009 mmole of hydroxylamine, in approximately 75 ml of solution, a t potentials 50-150 mv more positive than the half-wave potentials.
controlled-potential coulometry a t large stirred mercury-pool electrodes. The two sets of values are not in exact agreement because of the relatively large uncertainties that afflict estimates of diffusion coefficients and their relation to polarographic diffusion currents, but they are sufficiently concordant to demonstrate that the variations of IT'" are almost entirely due to changes in the average n value. The half-wave potential of the wave in 1 F sodium hydroxide, -0.390 v us. sce, is far more positive than the standard potentials of all of the couples involving hydroxylamine and its oxidation products. Although nitrous oxide usually predominates among the products obtained from the electrolytic oxidation of hydroxylamine, and although the n values listed in Table I show that it must also predominate under the conditions of these experiments, the standard potential of 5H20 4e = 2NH20H the half-reaction, N20 40H-, is -1.29 v os. sce13 (Stockholm convention). KO better agreement is obtained by postulating any other product. This is conclusive proof of the irreversibility of the process. Values of the parameter (1 - a)n, were obtained from plots of - E d e us. log i / ( i a- i) and from values of E*/,- Ell4as described by Meites and Israel,14 and also
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The Journal of Physical Chemiatry
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from the slopes of plots of log i us. E d e at potentials near the foot of the wave as suggested by Laitinen and Subcasky.15 I n 0.1-10 F sodium hydroxide, the value of (1 - a)n, is constant within experimental error and is equal to 1.70 0.02. This can hardly be reconciled with the assumption that n, = 1, while assuming that n, has any value above 2 makes it impossible to construct a credible explanation for the fact that the overall n value falls below 2 in the most strongly alkaline solutions investigated. We shall assume that n, = 2 in this range, which is consistent with earlier conclusions about the mechanism of the homogeneous oxidation of hydroxylamine as well as with the above value of (1 - a)n,. The log plots in these media were perfectly straight throughout the range of values of the log term from -0.8 to + O X In 0.1 F sodium hydroxide, the half-wave potential becomes more negative as the drop time increases; the expected dependence is dEll,/d log til, = -29.57/(1 - a)n, = -17.4 mv, and ten values obtained with drop times between 3 and-12 sec conformed to this prediction within a maximum deviation of only about 1.5 mv. At pH values below 13, however, the value of (1 a)n, is pH dependent, as is shown in Table 11. Its variation might be attributed to the doublelayer effect were it not for the fact that n also decreases over this range of pH values, as is illustrated by Table I and discussed further below. For reasons that will be explained in a later paragraph, we prefer to believe that the decrease of (1 - a)n, reflects a real change in the mechanism of the oxidation, that the rate-determining electron-transfer step involving two electrons a t very high pH values gives way to another involving only one at much lower ones, and that each of these processes occurs to some extent at pH values between 6 and 12. The fact that log plots are appreciably curved (both a t high and at low currents, so that their curvature cannot be ascribed to the erroneous selection of values of id) tends to support this belief. The values of (1 - a)n, recorded in Table I1 are averages over the rising portions of the waves. The effect of pH on the half-wave potential is also shown in Table 11, which includes values of the fundamental parameter ED1l2. This may be regarded16as the half-wave potential that would be obtained if t1l2>the drop time a t the half-wave potential, were 1 sec. A
(13) W. M. Latimer, "The Oxidation States of the Elements and their Potentials in Aqueous Solution," 2nd ed, Prentice-Hall, Inc., Englewood Cliffs, N. J., 1952. (14) L. Meites and Y . Israel, J . Am. Chem. SOC.,83, 4903 (1961). (15) H. A. Laitinen and W. J. Subcasky, ibid., 80, 2623 (1958). (16) Reference 7, pp 242-248.
ELECTROOXIDATION OF HYDROXYLAMINE
3623
Table 11: Polarographic Half-Wave Potentials and Values of (1 - a)na for the Oxidation of Hydroxylamine Supporting electrolyte
(1
0.05 F KHzPOr 0 , 0 5 F NanHP04, pH 6.52 0.02 F borax0.09 F KNOa, pH 9 . 1 8 0.01 F NaOH0.09 F KNOa 0 . 1 F NaOH 1.OF NaOH 5 . 0 F NaOH 10 F NaOH
-
a)no
EI/z, v us. ace
@V2*
aec
v us. ace
,..
..,
t1/2,
1.1
...
1.18=tO0.O1
-0.024
3.92
-0.009
1.45 f 0.05
-0.254
4.01
-0.242
1 . 6 9 i 0.09 1 . 7 0 2t 0.06 1.79 1k0.03 1 . 6 5 f 0.02
-0.332 -0.390 -0.428 -0.466
4.12 4.37 4.44 4.55
-0.321 -0.379 -0.417 -0.454
plot of EO1/,us. pH consists of two nearly linear segments. The slope of such a plot is given by
in a prior equilibrium, hydroxylamine is so very weakly acidic that these alternatives seem quite improbable. Bray, Simpson, and MacKenzie19 long ago concluded that nitroxyl was formed as an intermediate in the oxidation of hydroxylamine by iron(II1). Similarly, applying eq 1 to the data obtained in the range 9 2 pH I 13, using an average value of 1.3 for (1 - a)n,, gives p = 1.88 (ie., 2). The same end result is obtained from an argument based on the intercepts of plots of log i vs. E d c near the feet of the waves obtained at pH 6.5 and 9. Thus p = 2 over the whole range of pH values investigated, and since nu appears to approach 1 in the most acidic solutions, where hydroxylammonium ion is the predominating species, we may write “SOH+
+ 20H- +HNOH + 2Hz0 + e
as the rate-determining step in these media. The following mechanism serves to account for all of the experimental observations 2NOH + NzO
+ H20
+ HzO +H&’(OH)2 HN(OH)2 + 30H- +N02- + 2H20 + 2e NOH + n”20H + 2H20 X2 + 2H20 2HNOH + (HNOH)2 KOH
a t 2 5 ” ) where p is the number of hydroxyl ions consumed through the rate-determining step. Values of aoH-, the activity of hydroxyl ion in C F solutions of sodium hydroxide, were obtained by writing a O H - = f&, where fit is the corresponding mean ionic activity coefficient;” values of the pH were then obtained by log aoH-. At pH values above writing pH = 14 about 13, where (1 - a)n, is constant and equal to 1.70, the experimental value of dEol/,/d(pH) is -51 mv. According to eq 1, this corresponds to p = 1.48. Two sources of error can be discerned. One is that, according to Kielland’s calculations,’* the effective ionic radius of hydroxyl ion is smaller than that of sodium ion so that assuming them to be equal (as is tacitly done in the above procedure) leads to an overestimate of the pH, and the error increases as C increases. The other, and no doubt much the more important, is that the liquid-junction potential, whose sign is such that the half-wave potential is shifted toward more positive potentials, increases rapidly as C increases. It is not possible to make an exact correction for either of these effects, but because both of them tend to lower the apparent value of p , it seems impossible to escape the conclusion that p = 2. Accordingly, the rate-determining step at pH 2 13 may be written
+
NHzOH
+ 20H-
+NOH
+ 2Hz0 + 2e
(2)
Though this is kinetically indistinguishable from a scheme in which either ”OHor NOH2- is formed
(3)
--j
IYZ
(4) (5)
(6)
(7)
(8) If the pH between about 13 and 14.5, most of the nitroxyl that is formed in reaction 2 decomposes according to reaction 4. This path corresponds t o n = 2. A part of the nitroxyl, however, undergoes hydration a t a finite rate, yielding dihydroxylammonia (which has been assumed by earlier workers to be an intermediate in the reduction of nitrous acid13 and is therefore an attractive precursor of nitrite here), and the further oxidation of this to nitrite ion leads to an over-all n value that exceeds 2. That nitroxyl should have an appreciable half-life under these conditions is scarcely credible. Consequently, reaction 4 must be quite fast, and so also must reaction 5 because it would be indetectable if it were not. Reactions 2, 5 , and 6 thus constitute an ECE (electrochemical-chemical-electrochemical) sequence in which the intervening chemical step is much faster than the mass-transfer processes that control the rates of the electrolytic steps in a controlled-potential electrolysis on the plateau of the wave. KarpZ0has proved that a plot --j
(17) G.P. Haight, Jr., J . A m . Chem. SOC.,7 5 , 3848 (1953). (18) J. Kielland, i b i d , 59, 1675 (1937). (19) W. C. Bray, 14. E. Simpson, and A . A . MacKenzie, ibid., 41, 1363 (1919). (20) S. Ksrp, Ph D. Thesis, Polytechnic Institute of Brooklyn, 1967.
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GADDERAMACHANDRA RAOAND LOUIS~ I E I T E S
3624
of log i vs. 2 must be concave downward in such a case. The effect of a competing second-order side reaction that consumes the intermediate in an ECE process, as does reaction 4 here, has not been rigorously investigated, but a qualitative prediction is not difficult to make. The occurrence of reaction 4 tends to decrease the current that is obtained from reaction 6, and the decrease must be more pronounced near the beginning of the electrolysis than near its end because the pseudofirst-order reaction 5 must consume a larger and larger fraction of the nitroxyl as the electrolysis proceeds. The net effect will be an exaggeration of the downward concavity. Plots of log i us. t for controlled-potential electrolyses in 0.1 and 1 F sodium hydroxide were markedly concave downward after correction for the residual currents obtained with the supporting electrolytes alone. In still more strongly alkaline solutions, the value of n decreases below 2 and plots of log i us. 1 for controlledpotential electrolyses become perfectly linear after correction for the residual currents. These observations may be attributed to the occurrence of reaction 7, which would give n = 1 and a linear plot of log i LIS. t (corresponding to the over-all half-reaction, 2NHzOH 20H- + S, 4H20 2e, whose rate would be controlled by the rate of mass transfer of hydroxylamine to the electrode surface) if it were instantaneous. The increasing predominance of reaction 7 as the solutions become more alkaline is consistent with the results of Bray, Simpson, and ,IlacIlenzie.lg If the pH is below about 13, the value of n again decreasas below 2, for reaction 3 competes more and more effectively with reaction 2 as the pH decreases, and the decomposition of its product by reaction 8 would give n = 1. If the variation of (1 - a)n, in this range were attributed wholly to the double-layer effect, nitroxyl would have to be assumed to be the product of the rate-deterniining electron-transfer step here as well as in the more alkaline solutions. The fact that the rate of reaction 4 increases with increasing aciditylg would then give rise to grave difficulties in attempting to explain why n does not approach 2 as a limit as the pH becomes smaller. Voltaiiziiaetry at Platinum Electrodes. Though the electrooxidation of hydroxylamine at mercury electrodes thus appears to involve steps very similar to those previously postulated to explain its reactions with chemical oxidizing agents in homogeneous solutions, its behavior at rotating platinum wire electrodes is greatly complicated by variations in the nature of the eiectrode surface. A typical voltammogram obtained on backward polarization (i.e., toward increasingIy negative poten-
+
+
The Journal of Physical Chemistry
+
E,volts 08
YS
S C.E.
0.6
04
0.2
--2
co
a rn a
-I
--4!,
3
X'
Figure 1. The lower curve is the voltammogram of a solution containing 0.495 m F hydroxylamine in 0.10 F hydrochloric acid, obtained with a rotating platinum wire electrode by preanodization a t 1.2 v and polarization toward increasingly negative potentials. The upper curve is the residual-current curve obtained under identical conditions.
tials) in 0.1 F hydrochloric acid is shown in Figure 1. The residual-current curve shows the familiar cathodic peak, at about 0.45 v us. sce, due to reduction of the oxide film formed at more negative potentials. Because the height of this peak is not detectably affected by the presence of 0.5 m F hydroxylamine, the chemical reduction of the film must be much slower than the anodic regeneration of the film under these conditions. The wave is conspicuously asymmetrical : any reasonable extrapolation of the currents between 0.9 and 0.7 v indicates that the current should be extremely small around 0.6 v, but as the electrolytic reduction of the oxide filni begins, the anodic current resulting from the oxidation of hydroxylamine rises above the values thus extrapolated. Hence, the oxidation of hydroxylamine must be much more rapid on a finely divided platinum surface than on an oxide-covered one. Since the wave is always very near the cathodic peak or anodic wave resulting from oxide formation or reduction, it is also clear that the Oxidation is'much more rapid on the finely divided platinum formed by electrolytic or chemical reductionz1than it is on massive platinum. With 0.1 m F hydroxylaniine in a citrate buffer of pH 4.8, the height of the cathodic peak due to oxide-film reduction is only about half as large as in the absence of hydroxylamine, and it decreases further as the concentration of hydroxylamine rises until the peak cannot be discerned at all when the concentration of hydroxylamine is 0.6 mF. Hence, the rate of the chemical re(21) F. C . Anson, A n a l . Chem., 33, 934 (19G1).
3625
ELECTROOXIDATION OF HYDROXYLAMINE
duction of the oxide film increases as the pH and concentration of hydroxylamine increase; under the conditions just mentioned, it must exceed the rate a t which the film can be anodically regenerated, so that the rate of this regeneration exerts an important influence on the rate at which hydroxylamine is oxidized. Possibly because the plateau of the wave is longer in this medium than in a more strongly acidic one, a shallow minimum at about 0.6 v is obtained with 0.1 m F hydroxylamine on backward polarization, which gives E,,2 = 0.28 v. On increasing the concentration of hydroxylamine, the current a t 0.9 v increases more rapidly than does that at 0.4 v, so that with 0.6 mF hydroxylamine there appear to be two waves, the first of them very ill-defined. There is also a cathodic wave at -0.3 v, just preceding the final current rise. With 0.1 m F hydroxylamine, this is only about half as high as the anodic wave, and its relative height decreases further on increasing the concentration of hydroxylamine. As the cathodic wave is not obtained on forward polarization, it must be due to some product of the oxidation of hydroxylamine. The fact that the height of the cathodic wave increases only very slowly as the concentration of hydroxylamine rises about 0.1 m F could be explained in either of two ways. One is to assume that the substance responsible for it is adsorbed onto the electrode surface; the other is to ascribe the second anodic wave to the oxidation of nitrous acid22and assume that the formation of nitrous acid competes more and more effectively with that of the substance responsible for the cathodic wave as the concentration increases. Forward polarization in this supporting electrolyte gives a distinctly different curve, as is illustrated by Figure 2. Not only is the plateau much more distorted and the half-wave potential much more positive (0.71 v as against 0.28 v) than on the curves resulting
Figure 2. The lower curve is the voltammogram of a solution containing 0.0989 mF hydroxylamine in a 0.1 F sodium citrate buffer, pH 4.78, obtained with a rotating platinum wire electrode by precathodization a t -0.45 v and polarization toward increasingly positive potentials. The upper curve is the residual-current curve obtained under identical conditions.
from backward polarization, because of the difference between the potential a t which the oxide film is formed and that at which it is reduced, but in addition, the wave height is roughly twice that obtained on backward polarization (which in turn is roughly twice that obtained on backward polarization in 0.1 F hydrochloric acid). These observations prove that the oxide film and the finely divided platinum formed by its reduction are intimately involved in the oxidation of hydroxylamine on platinum electrodes. Far-reaching modification of the mechanism that suffices to describe the behavior on mercury would be needed in attempting to account for these phenomena and those previously reported by others. -5 ~~
~
~
~
~_____
(22) R. N. Adams, “Handbook of Analytical Chemistry,” L. Meites, Ed., McGraw-Hill Book Co., Inc., New York, N. Y . , 1963, Section 5 , p 151.
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