NOTES
1126
we can calculate the heat of formation of ethyl bromide from the reaction Hz v CHs-CHoBr CH2=CHBr AH = -34,570 cal./mole The value of -15.9 kcal./mole compares with the value of -13.0 given in reference 4. Linnett? and co-workers have studied the equilibrium between ethylene, hydrogen bromide and ethyl bromide. From their data, they calculated the heat of formation of ethyl bromide to be -15.3 kcsl. at 25’. The addition of hydrogen to vinyl bromide is apparently stepwise and involves, first, substitution of the bromine and then saturation of the double bond in ethylene. Since the heat of hydrogenation of ethylene is known,8 the heat involved in substituting the bromine may be calculated. The results are given in Table 11. The value of -15,370 cal./mole is about 3000 cal. less exothermic than the corrosponding reaction of vinyl chloride.6
+
(7) M. R. Lane, J. W. Linnett and H. G. Oswin, Proc. Roy. Soc. (London), A216, 361 (1953). (8) G . B. Kistiakowsky, H. Romeyn, J. R. Ruhoff, H. A. Smith and W. E. V a u g h n , J . Am. Chem. Soc., 61, 65 (1935).
POLAROGRAPHY OF THE FERRICYANIDE ION IN THE ABSENCE OF I N E R T ELECTROLYTE BYP. KIVALO Finland Institute of Technology, Department of Chemistry, Helsinki, Finland Received April 1 1 , 1967
The polarographic reduction of ferricyanide produces a well-defined diffusion current which is directly proportional to the concentration in various supporting electrolytes, e.g., 0.1 M KC1 or NaC104.I The reduction starts a t a positive potential referred to S.C.E. and is accompanied with
,1
0
-0.5
-1.0
- 1.5
- 2.0
E , v. us. S.C.E. Fig. 1.-Polarograms of 1, 5 and 8 mhf concentrations of ferricyanide obtained without supporting electrolyte and using a Leeds & Northrup Electrochemograph recording polarograph. The corresponding cell resistances are 8500, 2700 and 2000 ohms; t1/8rn*‘8 = 1.84. 6 -
a4-
d ’
+-
Vol. 61
a pronounced maximum which, however, is easily suppressed. Frumkin and Florianovich2 have recently reported to have found a “dip” in the diffusion plateau of the ferricyanide ion. The current is said to drop between -0.4 and -0.8 v. us, S.C.E. for concentrations below 0.01 N and in the absence of inert electrolyte. The authors were able to observe only the fall in the current without a subsequent rise in it. Since the polarographic reduction of ferricyanide is of particular interest from the point of view of a proposed reduction mechanism of certain other anions3 that produce minima in their diffusion plateaus (which, however, also show a subsequent rise in the current to the original diffusion value), its reinvestigation was deemed desirable. In Fig. 1 polarograms of 1, 5 and 8 m M solutions of ferricyanide obtained in the absence of supporting electrolyte are shown. It is clearly seen that a very pronounced maximum starts the waves and that the maximum drops off between -0.5 and -0.6 v. us S.C.E. Furthermore, the limiting currents are practically constant. The diffusion control of the current was tested using 1 mM concentration at -1.0 v. varying the height of the mercury column h,
om.
drn i/io
40
50
60
70
80
1 1
1.12 1.11
1.23 1.21
1.32 1.29
1.41 1.36
It may be mentioned that the maximum was present even a t as low a concentration as 0.2 mmole per liter. I n the absence of inert electrolyte, the limiting current is lower than the “true” diffusion current owing to the migration current which in this case decreases the latter.4 Addition of supporting electrolyte should increase the limiting current. The results of these experiments are presented in Fig. 2. The concentration of the ferricyanide was 1 mmole per liter. Upon increasing the NaC104 concentration from 0 to 0.05 Ai! the plateau current rose to a value which then remained practically constant and independent of further addition of the supporting electrolyte. At the same time it is seen that the maximum is suppressed and disappears above 0.05 M concentration of NaClOa. An increase in the ferricyanide concentration again develops the maximum. The diffusion current constant calculated from polarogram 4 in Fig. 2 gives a value of 1.84 in good agreement with the value reported by M e i t e ~ . ~ The relation iol,/& was here found to be 0.7 which agrees with the calculated value based on io,/& = 1/1 Tu- where To- is the anion trnnsference number, assumed t o be equal to 0.L6 Based on the above findings the present author
+
20
- 1.0 E,v. us. S.C.E.
-2.0
Fig. 2.-Polarograms of 1 mM ferricyanide solutions obtained with a Radiometer Model P03k recording polarograph a t various concentrations of the supporting electrolyte; t1/srn*’a = 2.28: Curve 1, 0 M NaC104; Curve 2 , 0.005 M ; Curve 3,0.04 M ; Curve 4, 0.05 M . (1) J. J. Lingane and I. (1939).
M . Kolthoff, J. Am. Chsm. Soc., 61, 825
(2) A. N. Frornkin and G . M. Florianovich, Doklady Akad. Nauk, S.S.S.R., 80, 907 (1951). ( 3 ) P. Kivalo and H. A . Laitinen, J . -am. Chem. Sac., 77, 5205
(1955). (4) I. 11. Kolthoff and J . J. Lingane, “Polaroqraphy,” Vol. I , 2nd Ed., Interscience Publishrrq. Inc., New York. N. Y., 1952, P. 122. (5) L. Meites, J . A m . Chem. Soc., 7 3 , 1581 (1051). (6) B. E. Conway, “Electrochemical Data,” Elsevier Publishing Co., New York, N. Y., 1952, p. 166.
0
r
b
I
NOTES
August, 1957
1127
proposes for the “minimum” of the ferricyanide found by Frumkin and Florianovich an alternate explanation according to which it simply is a polarographic maximum on its way out. The sensitiveness toward the presence of inert electrolyte, as reported by the above authors, can very likely be attributed to the decrease of the negative migration current. In concluding it is worth emphasizing that the limiting current plateau of this triply charged anion also in the absence of inert electrolyte is flat. Thus the increasingly negative made surface charge of the electrode does not seem to be a hindrance for the diffusion of the highly negative ion to the electrode, However, the electrostatic effect due to the migration is clearly observable.
ANOMALOUS EFFECT OBSERVED I N T H E STUDY OF INTERMOLECULAR HYDROGEN BONDING OF METHANOL I N A NONPOLAR SOLVENT BY KENNETH M. SANCIER~ COntdbUtiOn from The Stanford Research Institute, Menlo Park, California, and The Johns Hopkins University, Chemzetry Department, Baltzmore, Maryland Received April 88, 1867
It has been shown by infrared ~ p e c t r a , ~by -~ heat of mixing, and by orientation polarization6 that intermolecular hydrogen bonding occurs a t concentrations less than 0.01, and probably as low as 0.001 mole fraction methanol in carbon tetrachloride solution. The evidence indicates that the intermolecular species present in dilute solutions are principally the trimer and, a t extreme dilution, the dimer. Vapor-liquid equilibria of alcohols in various non-polar solvents have been studied,6 ~ 7 but no information concerning the nature of the intermolecular species present was deduced. I n the present study, two different methods of analysis of vapor-liquid equilibria were used to examine systems of methanol in inert solvents down to concentrations a t which virtually all intermolecular hydrogen bonding should be eliminated. The objective was to determine the extent to which the dissociation energy of a hydrogen bond is different for the dimer and trimer a t higher concentrations. This objective was not realized because adsorption of methanol on glass a t very low concentrations becomes important enough to mask hydrogen bonding effects. However, the experiments are of interest because they demonstrate a source of serious error which may be encountered in glass systems, and because the methods of analysis may be valuable if applied in an all-metal system. (1) The Stanford Research Institute, Menlo Park, Cal. (2) J. J. Fox and A. E. Martin, Proc. Roy. Soc. (London),A162,419 (1937). (3) J. Errera, R. Gaspart and H.Sack, J . Chem. Phys., 8 , 63 (1940). (4) F. A. Smith and E. C. Creitz, J . Research Natl. Bur. Standards, 46, No. 2, 145 (1951),Research Paper 2185. ( 5 ) K. L. Wolf, Trans. Faraday SOC..33, 179 (1937). (6) (a) G. Scatchard, 8. E. Wood and J. M. Mochel, J . A m . Chem. SOC.,68, 1960 (1946); (b) 8.E. Wood, ibid., 68, 1963 (1946). (7) A. Niini, Ann. Acad. 8ci. Fe’snnffiae,ASS, No. 8, 1 (1940).
2 - L O P N,
I
.
3
4
Fig. 1. Method. 1.-Vapor-solution equilibria of the methanolcarbon tetrachloride system were studied* in the concentration range from 1.0 to 0.0002 mole fraction ?ethanol and a t five temperatures ranging from 10 to 50 . After drying and deaerating, the solution and its vapor were equilibrated in a glass, high-vacuum system with mercury manometers and Hoke metal sylphon bellows valves, and thermostated to &0.002°. Means had been provided to obtain the following data: total pressure P and mole fraction methanol in the Sam les of vapor and solution, naand N,, respectively. Mole gaction methanol was determined from the total pressure of a given gas volume of the sample condensed in a small, fixed volume, a t a fixed temperature, in an apparatus which had been calibrated previously with samples of known compositions.g Raoult’s law was applied to these data to calculate P.*,the hypothetical vapor pressure that “pure” methanol would exert if the extent and nature of hydrogen bonding were not altered - =
N.
P.*
In Fig. 1 are plotted the values of log P.*against -log N , for five temperatures. The value of P.*steadily increases 0.004),a t to a maximum at about -log N . = 2.5(Na which point its value unexpectedly decreases. At such dilution P: would be expected not to decrease but rather to approach a constant high value, with the methanol behaving like an ideal liquid (or gas) with no association. A corresponding anomaly was observed when the apparent values of A H v , the heat of vaporization of methanol, were plotted against -log N.. These values were calculated in the usual manner from the slopes of the log P.*against l/T°K. curves. The value of A H v decreased regularly from 9 to 5.2 kcal./mole of methanol as the mole fraction of methanol was decreased from 1.O to 0.001. It then unexpectedly increased to about 9 kcal./mole methanol when N , became (8) Baaed on a Thesis submitted by Kenneth M. Sancier to the Johna Hopkina University, October, 1949, in partial fulfillment of the requirements for a Ph.D. degree. (9) K.M. Sancier, Anal. Chsm., 44, 1668 (1962).
