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 to a maximum at about -log N . = 2.5(Na 0.004),a t 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).
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less than 0.001. It would be expected that the contribution of hydrogen bonding to the heat of vaporization of methanol would become smaller and eventually negligible as the methanol concentration is decreased, and that the heat of vaporization of methanol would not increase again but rather approach a low constant value. Method 11.-The behavior of the equilibria systems in the foregoing experiments was most unusual, and it appeared worthwhile subse uently to check the results. A more sensitive method analysis employing C14-taggedmethanol was usedl0 to extend the measurements into a more dilute concentration region. Cyclohexane was substituted for carbon tetrachloride because of potential difficulties in the purification of the latter. A similar glass, high-vacuum system was employed. The solvents, Eastman Kodak Spect,roGrade stored in glass flasks attached to the vacuum system, were dried in situ with calcium hydride and were thoroughly deaerated. The analytical procedure employed an internal counter Geiger-Muller tube. For analysis of a as sample, the GM tube was filled with 10 mm. of t8e organic sample (methanol, cyclohexane or mixtures of these two behaved equallywell) and then 90 mm. of argon. The GM tube had a 1% lateau over a range of 200 volts and a starting potential oPabout 1200 volts. Only lo-' mole of sample is required for this analysis, compared with 15 X mole for method I. With the specific activity of the C14HaOH (Tracerlab, 1 mc./mmole) known, and with 100% counting efficiency assumed, a methanol concentration of 1O-B mole fraction theoretically could be detected with a counting rate of 1000 c.p.m. A serious source of error was detected immediately. The memory of the GM tube after counting a sample was large, of the order of 95%. This memory effect was eliminated when the GM tube was washed with liquid methanol and dried at 100". However, it soon became evident that considerable adsorption was occurring on the glass walls of the high-vacuum system used to transfer the gas samples. These adsorption effects appeared to be extensive enough to invalidate the intended measurements. Subsequent work was therefore devoted to semi-quantitative evaluation of the extent of adsorption in the glass system in order to determine whether the method was applicable. Standard gas-counting mixtures were made up, each of which contained the same concentration of Cl4H8OH, and sufficient untagged methanol was then added to provide gas mixtures containing 10-2 and 10-1 mole fraction total methanol in cyclohexane. Adsorption of these mixtures in a 5-liter flask was permitted to proceed for 24 hours. The mixtures were then individually introduced into the GM tube in a standard way. The counting rates, corrected for background, were materially less with lower methanol concentrations (for the same original C14HaOH concentration): 8, 32, 70, and 85 c.p.m. for IO-', 10-8, and 10-1 mole fraction methanol, respectively. The highest counting rate observed for the original sample of solution (made up to be 1000 c. .m.) was 435 c.p.m. This was obtained by pretreatment o r t h e glass walls of the connecting tubing through which the sample was transferred to the GM tube with 5 mm. of methanol. Further evidence of adsorption of methanol by glass was obtained when the treatment of a sample in the connecting tubing was varied. It is estimated that approximately a monolayer of methanol is adsorbed from a gas containing 10-2 mole fraction methanol onto a glass surface, with a perfectly smooth surface assumed; about 1/10 less is adsorbed from lo-' mole fraction methanol. The counting efficiency of the GM tube was not materially altered by the addition of tracer-free methanol after the sample was introduced. However, if tracer-free methanol is mixed with the active sample prior to introduction in the GM tQbe, the counting rate is nearly doubled. The indication is that the effective geometry has been doubled in the latter case by the reduction of the adsorption on the walls of the GM tube. It should be pointed out that the counting rate data on adsorption of methanol on glass are only qualitative, for no true equilibrium was attained. The data are comparable only because the handling of all samples was carefully duplicated.
07
(10) This work was done at and supported by Stanford Research Institute, Menlo Park, California.
Vol. 61
Summary of Results Experimental evaluation of intermolecular hydrogen bonding of methanol was not feasible, particularly in dilute solutions of methanol, because of adsorption of the methanol on glass. This adsorption appears to be not unlike that of water on glass. Equilibria studies of the vapor-solution type in glass containers therefore seem impractical. Not only is it difficult to transfer samples from the equilibration chamber to the analysis section, but adsorption in the glass equilibration chamber itself competes with intermolecular hydrogen bonding for methanol. The data obtained by the method employing the C14 tracer provide an explanation for the previous anomalous results (method I). The strong adsorptive effects, active a t concentrations less than 0.1 mole fraction methanol, and very pronounced a t concentrations of about lou3 mole fraction, could account for the exceptional decrease in partial pressure of the methanol (Fig. 1) and increase in heat of vaporization at concentrations less than mole fraction methanol. Once the problem of adsorption is minimized (possibly by use of an all-metal system), the two methods described above may be suitable for the study of the nature of hydrogen bonding in very dilute solution. The high sensitivity of the tracer method suggests that it might be used for the evaluation of adsorption phenomena.
A METHOD FOR THE RAPID DETERMINATION O F THE SOLUBILITY OF GASES I N LIQUIDS AT VARIOUS TEMPERATURES BY FRANK J. LOPREST Chemistry Department, Reaction Motors, Inc., Denvills, New Jsraay Received May 3,I967
A need for a large amount of gas-liquid solubility data which are not available in the literature arose in our laboratory. The solubility data were required over a range of temperatures and a precision and accuracy of = k l % were considered adequate for the application of the data. It is believed that the simple apparatus which was developed to fulfill our need may be of general interest. Methods for the determination of the solubility of gases in liquids have been critically reviewed.lP2 The recent work of Cook2in which the solubility of hydrogen in certain non-polar solvents was determined contains a particularly complete discussion of the problems attending these solubility measurements as well as a discussion of previous work in this field. Cook sets the specification on his measurements that they be reproducible to *0.05%. He describes a rather elaborate apparatus with which he was able to obtain the specified reproducibility. Jn view of our requirements a much simpler apparatus and method than those used by Cook were developed and these are described in this paper. Apparatus and Procedure.-The method consists of the introduction of a measured quantity of gas into a constant
A. E. Markham and K. A. Kobe, Chem. Reus., 28, 519 (1941). (2)-M. W. Cook, U. 8. Atomic Energy Comm., UCRL-2459, 1954. (1)
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