LEONARD NEWMAN,
1814
JOAO DE
0. CABRAL AND DAVIDN. HVME
VOI.
80
effects of this order of magnitude.8 The N P + ~ tion of the dismutation products of the activated (aq.) ion presumably is formed by further reac- complex with hydrogen ions. (8) Cj.. V. Gold and D. P. N. Satchell, @rrart. Reo.. 9, 51 (1955).
(COSTRIBUTION FROM
THE
LEMONT, ILLINOIS
DEPARTMENT OF CHEMISTRY AND LABORATORY FOR NUCLEAR SCIENCEO P
THE
MASSACIIUSETTS
INSTITUTE O F TECHNOLOGY]
A Polarographic Study of Mercuric Cyanide and the Stability of Cyanomercuriate Ions' BY LEOSARD N E W M A N , JOAO DE 0 . CABRAL
AND
DAVIDN. HUME
RECEIVED OCTOBER 28, 1957 The reduction of mercury cyanide and the oxidation of mercury in hydrogen cyanide were shown to be polarographically irreversible in acid solutions, becoming reversible in basic solutions. In basic solutions the cyanide released by the reduction can complex mercuric cyanide a t or around the surface of the drop. The species being reduced are Hg(CN)*, Hg(CN)?The logarithms of th: over-all formation constants of these species are 33.9, 38.1 and 40.6, a t an ionic and Hg(CN)a--. strength of 2.0 and a temperature of 30
.
Introduction The reduction of mercuric cyanide a t the dropping mercury electrode was first extensively studied by TomeY who found in the range pH 6.0 to 9.5 a reversible wave corresponding to the over-all electrode reaction Hg(CN)*
+ 2H+ + 2e-
--f
Hg
+ 2HCN
I n basic media the reaction became Hg(CN)2
+ 2e- -+ Hg 4-2CX-
The half-wave potentials in the pH range above 9.5 were observed to deviate from predicted values, and this was attributed to complexation of mercuric cyanide by cyanide ions liberated in the electrode reaction. Tome3 offered no experimental evidence for the formation of higher complexes, and the deviations were in the opposite direction from that expected to be caused by such an effect. It has been the purpose of the present investigation to study the polarographic behavior of mercuric cyanide over a much wider range of PH values and determine the nature and effect of possible complex formation. Since the completion of this study,' Tanaka and hlurayama3have reported an extensive study of the anodic wave of mercury in cyanide medium which corroborates our findings regarding the formation of higher cyanide complexes of mercury a t the drop and leads to values for the formation constants which are in substantial agreement with those determined froni our cathodic waves of mercuric cyanide. About the same time, Anderegg4 published the results of a potentiometric determination of mercuric cyanide stability constants. Theory Tome3 derived for the equation of the wave, the expression
and pointed out that E l / , was not independent of concentration. I n place of El/, it is more conven(1) Taken in part from the Doctoral Dissertation of Leonard Newman, Massachusetts Institute of Technology, 1956. ( 2 ) J. Tome:, Collection C5cch. Chem. Commtrn., 9, 81 (1937). (3) N. Tanaka and T. Murayama, 2 . physik. Chem. N . F., 11, 366 (1957). (4) G . Anderegg, Helv. Chim. Acta, 40, 1022 (I'J.77).
ient to use the concentration-independent potential on the wave, which we shall call Ei and which leads to Ede = Et
+ 0.030 log ( '9)
for the equation of the wave, It is shown readily that KA COH . f ~f f m - ( 3 ) Et = E' 0.060 log
+
+
+
KA
in which E' = Eo
d2 + 0.030 log 1 dnKzf2cN-
(4)
where K A is the ionization constant of hydrogen cyanide, COH+is the concentration of hydrogen ion a t the electrode surface, f H + and f C N - are the activity coefficients of hydrogen and cyanide ions, Eo is the standard potential of the mercury-mercuric ion couple, dl and d, are the diffusion coefficients of cyanide and mercuric cyanide, respectively, and Kz the formation constant of mercuric cyanide. It is seen that Ede is equal to Ef when (id i)/i2 is equal to unity and that Er is independent of mercuric cyanide concentration. It should be noted, however, that the relative position of Ef on the wave shifts with concentration. If all measurements are done a t a fixed, high ionic strength, activity coefficients may be assumed to be constant. The equation of the wave predicts a plot of Ef vs. pH to be linear with a slope of 60 mv. per #H as long as CoH+rH+fCN- >> KA. For values of C'H'fH'fCN> pH+ + fH fCN-,we have, analogous to the sun1 of ecluations 3and4
+
Et
+ Eo f 0.030 1% d n R ndl"f n o v -
(61
where n is the cyanide-mercuric ratio of the particular complex. From Table I11 it can be seen that d, is independent of whether the reduction depends upon Hg(CN)Z, Hg(CN)Z- or Hg(CN14--. The average values of d, and dl calculated from idcathodic /C'H~(CN)* and i d a n o d i c ' C ' c N - are 8.2 X loa and 7.1 X lo3, respectively. With EOfor the mercuric-mercury couple equal to 612 mv. vs
LEONARD NEWMAN, JOAO
1815
Worker
This work Brigando and Joh’ Sherrillls Charlot and Gauguin’? Gauguin14 Tanaka and M u r a y w ” Anderegg4
DE
0.CABRAL AND DAVID N. HUME
Log Ka
Log Kr
Log Kd
Log k3
33.9
38.1
40.6
4.2 3.9
..
..
..
21.7 35.3 34.7
.. .. , .
.. 39.0 38.6
.. 41.4 40.5 27.7 41.5 41.5
.. .. ..
3.7 3.8
Log fir 1,og ( K c / K d
2.5 2.9
6.7 6.8
.. ..
..
..
2.5 3.0
..
6.0 6.2 fi.8
\Tal.
p
Method
2.0 0.01 t o 0 . 0 2 0.05 to 0.2’
Polar. Poten. at 2’ Poten.
0 0.2 0.1
Cond. Polar. Poten.
. . . . ...
so
...
saturated calomel electrode and the Ef values -291, -300 and -260 mv. for Hg(CN)*, Hg(CN)3- and Hg(CN)d--, the over-all formation constants K,, K 3 and Kd can be calculated from equation 5 where
crepancy ; however, since Sherrill did not consider the formation of Hg(CN)Z-, his value for Kd should be too high. The values of K z and Kk as obtained from conductivity measurements, and reported by Gauguin, l4 are obviously too low; however, the value of K4/K2 falls in line with the one calculated in this study and with the one calculated from the study of Brigando and Job. It appears that we are the first to have determined values for K2 and K3 independent of and values reported for Kd and/or K A . The main difference between the values of our constants and those of Tanaka and Murayama if we assume that the activity coefficient of CN- is arise from the assumptions made in calculating K2. not only constant due to the high and constant Tanaka and Murayama assume that fCN- can be calculated from Kielland’s’j data a t ionic strength ionic strength, but is also equal to one. The stepwise formation constants k3 and k4 of 0.2, that glass electrode measurement of pH a t can be calculated from K3/Kz and K,/K3. A sum- p = 0.2 gives the activity of hydrogen ion directly mary of the logarithms of the formation constants and that K A is 4.0 X 10-lo. For our calculation of calculated in our study and those reported in the K2 we assume only that the fcx-a t p = 2.0 is equal to one. Whereas our values of K3 and K , are calliterature are included in Table IV. Brigando and Job,’ in a potentiometric study a t culated directly from the experimental data, 2’ and p = 0.01 to 0.02, reported formation con- Tanaka and 11Iurayama utilized Sherrill’s13 data for Kd and their value of Kf in the calculation of stants of K S . The results of Anderegg, calculated by application of the Bjerrum average ligand number technique to potentiometric data obtained with the and glass and mercury electrodes are in good agreement with ours and those of Tanaka and Murayama, The value of K B and Kd and Kd/Kz can be cal- especially, in view of the differences in experimenapproach. culated from the ionization constant of hydrogen talTo further demonstrate that the value obtained cyanide. The value of this quantity is not known in this study for the Ef of Hg(CN)4-- is valid, accurately : various experimental determinations in polarograms run on mercuric cyanide with a the literature have given values of 7.2 X 10-10,4p8 large excess ofwere cyanide in basic solution. It can be 13 X 10-10,9and 260 X 10-lo.lo Latimerll gives assumed that all the mercuric cyanide is coriiplexed the value 4 X calculated indirectly from as Hg(CN)4--. For a inixecl anodic--cathodic thermodynamic data, as the most probable one, and wave we have used this value to derive the additional constants from the data of Brigando and Job. Considering the differences in temperature and ionic strength, the agreement between their con- where i d c and ida represent the anodic and cathodic stants and ours is quite good. The value of IG diffusion currents. When cyanide is a lot greater which we have obtained also agrees well with the than mercuric cyanide, the I d a becomes a lot greater value reported in the text by Charlot and Gau- than i; and the denominator becomes a constant. guin,12but only fairly well with the potentiometric Therefore value reported by Sherrill.13 The differences in &e = E,, C X 0.030 log ( h a - i) (7) ionic strength might account for some of this diswhere in basic solution
+
( 7 ) J. Brigando and P. Job, C o m p l . r e n d . , 222, l 2 ! l i ( l 9 - l G ) . (8) H. Lunden, Z. p h y s i k . C h e m . , 64, 532 (19lIG), (9) J. Walker and W. Cormack, J . Chem. SOC.,77, 5 (1900). (10) J. L. R. Morgan, 2. physik. Chem.. 17, 513 ( I S M ) . (11) W. h l . Latimer, “Oxidation Potentials,” 2 n d Ed., PrenticcHall, N e w York, N. Y . , 19.52, p. 137. (12) G. Charlot and R. Gauguin, “Les Methodes D’Analyse des Reactions en Solution,” Masson e t Cie, Editeurs, Paris, 1951, p. 315. (13) M. S. Sherrill, Z. p h y s i k . Chem., 43, 705 (1903); 47, 103 (1004).
Ef,CN = Eo - 0.030 log d4k’4(CN-)4S4CN- (8)
Equation 7 predicts that a plot of Ede vs. log ( i d c - i) should have a slope s equal to 30 mv. Experiment 19 of Table I11 shows that the proper slope for the postulated Hg(CN)4-- is attained. (14) R. Gauguin, Anal. C k i m . A d a , 3 , 489 (1949) (1.5) J. K ~ e l l m dTIIIS , J O U R N A L , 59, 1675 (1037)
-1pril 20, l%S
DECOMPOSITION OF BIPIIOSPHINE IN LIQUIDAarhion-IA
From E/, C N - = -502 mv., Eo = 612 mv., d4 = 8.2 X lo3, Ha./M, cyanide = 0.1 M and fcN- = 1.0, K4 was calculated from equation 8 as 1040.1. This can be considered good agreement with the K4 = 1040.6, obtained when mercuric cyanide was assumed to be complexed as Hg(CN)4-- a t the surface of the drop. It was thought necessary to test the polarograms for polarographic reversibility since the arguments above are based upon this condition. I n order to do this, anodic waves of cyanide were obtained in basic solutions. As can be seen in Table 111, the analyses of the anodic waves were the same as analogous cathodic waves. The sy5tem can therefore be said to be polarographically reversible. Another usually important criterion €or reversibility is the attainment of a single continuous polarogram for a mixed anodic and cathodic wave. It can be shown that with the type of electrode phenomena that we have observed this behavior is not to be expected. In an equimolar solution of cyanide and mercuric cyanide most of the cyanide is complexed by the mercuric cyanide. The cathodic portion of the polarogram should start with a dependency on the re-
[COSTRIBUTIOS PROM THE JOHN
1819
duction of Hg(CN)3-, with the reduction of Hg(CN)4-- taking over quite early. On the other hand, since the activity of the mercuric ion is decreased by the presence of mercuric cyanide, the anodic portion should start with a dependency upon Hg(CN)3-, with the reduction of mercuric cyanide taking over quite early. The cathodic portion would be mainly dependent upon the reduction of Hg(CN)d-- and the anodic portion upon the oxidation of mercury to mercuric cyanide. One should not expect to obtain a straight line over the anodic and cathodic portions from any plot of E d e vs. log ( i d - ;)/is. When this experiment was tried, the polarogram looked like a single continuous wave, but, as expected, no log plot gave a straight line with the proper slope over the anodic-cathodic portion. Acknowledgments.-This work was supported, in part, by the United States Atomic Energy Commission. J. deO. Cabral was the holder of a Massachusetts Institute of Technology Foreign Students Summer Program Fellowship in the summer of 1053, while on leave from the University of Oporto, Oporto, Portugal. CAMBRIDGE. MASSACHUSETTS
HARRISON LABORATORY O F CHEXISTRY,
UNIVERSITY OF PESNSYLVANIA]
The Lower Hydrides of Phosphorus. 11. The Decomposition of Biphosphine in Liquid Ammonia's2 BY EVANH. STREET, JR., DAVIDM. GARDNER AND E. CHARLES EVERS RECEIVED OCTOBER 22, 1957 Biphosphine decomposes iri liquid ammonia to lose phosphine, leaving a black residue of variable composition contaiiii,.g phosphorus, hydrogen and solvent. Products are obtained whose P/H ratio varies from 3.05 to 4.81 and whose NHI!H ratio varies from 0.09 to 0.92. Similar substances are formed on treating phosphorus with ethylamine. The c o i i \ t i t u t ~ i ~ i of these substances is discussed in relation to solid hydrides obtained when biphosphine is decomposed in ' J U C ~ L Oatid iu the presence of moisture.
I n our previous communication, we described the decomposition of biphosphine a t room temperature under anhydrous conditions. Decomposition in vacuo proceeded with the elimination of phosphine, forming orange-yellow solids whose ultimate compositions appear to lie in the neighborhood of P2.25H, (P9H4). With moisture present the products contained larger proportions of hydrogen ; materials obtained in a variety of ways have reported compositions averaging around PzH.3-5 These preparations left small amounts of nonvolatile residue on strong heating in vacuo. Another class of hydrides has been prepared by (1) Taken in part from a Thesis by Evan H. Street, J r , , presented in partial fulfillment of the requirements for the Ph.D. degree, June, 1955. (2) This research was supported in part by the Office of Naval Research under Contracts Nonr-598(00) and N8onr-74200. Reproduction in whole or in part is permitted for any purpose of the United States Government. Presented before the Inorganic Division of The American Chemical Society, 132nd National Meeting, New York, September, 1957. (3) E. C. Evers and E. H . Street, Jr., T H I SJOURNAL, 78, 5726 (1956). (4) (a) A. Stock, W. Bottcher and W. Lenger, Ber., 41, 2839, 2847, 2853 (1905); (b) R. Schenck and E. Buck, ibid., 37, 915 (1904). (5) The molecular formula PUHBwas proposed by R. Schenck a n d E. Buck, ibid., 37, 915 (1904), as a result of cryoscopic measurements in molten phosphorus.
treating solid hydrides6s7or white phosphorus6with a variety of amine solvents. Stock4" found that phosphine was eliminated and reddish colored solutions were formed when P2H was treated with liquid ammonia. The substance P4.5H, (PgHJ, which was reportedly obtained by the thermal decomposition of PzH under the proper cond.ition~,~" likewise dissolved, but without the evolution of phosphine. In all cases black substances were obtained by evaporating the solvent, whose eompositions averaged around P4.5H and which contained up to one molecule of ammonia. A small amount of non-volatile residue remained on strong heating in vacuo. This appeared to be P&J5.6a Products obtained by the direct action of liquid ammonia on phosphorus were similar. Although comprehensive quantitative data have been reported only in the case of products obtained from ammonia, it seems likely that other nitrogen bases also yield similar products. In view of the thermal instability of biphosphine (6) R. Schenck, ibid., 36, 979, 4202 (1903). (7) (a) A. Stock, ibid., S6, 1120 (1903); (b) A. Stock and 0. Johannsen, ibid.. 41, 1593 (1908); (c) H. Krebs, Z. anorg. Chein., 266, 175 (lY.51); Angew. Chenr., 64, 293 (1952).
