DISSOCIATION PRESSURES OF SOME DEUTERATES
Jan., 193s
87
tively low concentration of ozone is a bimolecular reaction with a specific rate of 50. (The units are moles per liter and minutes.) It is concluded that ozone and hydrogen peroxide react competitively with an intermediate substance formed in this bimolecular reaction, and that the ratio of the specific rates is 2.6 to 1. Two mechanisms are considered, each of which explains the results satisfactorily. They involve H02. intermediate formation of H20a or of HO Attention is called to the quantitative results of Harcourt, published in 1862, on the reaction of potassium superoxide, KOz, with dilute sulfuric acid and with water vapor. It is concluded that summary hydrogen peroxide is not catalytically decomposed Analysis of the rate data of Rothmund and in the presence of HOBand 0 2 - , and that the speBurgstaller on the interaction of ozone and hydro- cific rate of the reaction 2H02 = H202(aq) 0 2 , gen peroxide a t 0' in 0.01 N sulfuric acid shows is much greater than that of either 2HO2 = O3 HaO(l)or H02 HzOa(aq) = HO H20(1) agreement with the relation -d(Oa)/dt + -dThe decreases of free energy in the various inter(H8O2)/dt = 1 5.2(O3)/(H202). Their conclusions are confirmed: hydrogen peroxide is a mediate reactions are estimated. The conclusions presented in this paper differ catalyst for the decomposition of ozone, while ozone is not a catalyst for the decomposition of from those published by Weiss in 1935. hydrogen peroxide. The limiting result a t rela- BERKELEY, CALIF. RECEIVED SEPTEMBER 20, 1937
still sufficient uncertainty to warrant serious consideration of the simpler mechanism involving H a s formation in reaction (loa). If the mechanism does involve H203, reactions (loa), (lob) and (lob') should be one-directional to give the observed rate laws. Accordingly it is estimated that the free energy decrease in reaction (loa), HdMaq) Odg) = HzOdaq) Odg), lies between 17 and 47 kcal. Even with the lower estimate, H203(aq)is stable with respect to HO HOz. By means of this lower value, AF&8.1 for Ha03(aq) is calculated to be -270, and AH: for H203(g)is estimated to be -295 kcal.
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[CONTRIBUTION FROM THE
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FRICK CHEMICAL LABORATORY OF PRINCETON UNIVE!RSITY]
Dissociation Pressures of Deuterates of Cupric Sulfate and of Strontium Chloride' BY FRANCIS T. MILESAND ALANW. C. MENZIES The current interest in this subject is evidenced by the fact that, since the present work was completed, dissociation pressure measurements upon one of these deuterates have been reported by investigators in three other laboratorie~.~-5 Materials and Method.-The materials were from the same samples as those described elsewhere.6 An additional check upon the concentration of the deuterium water was made possible by a melting point determination on a portion distilled from the deuterates t o another part of the sealed apparatus employed. The method used was a comparative one in which the difference of dissociation pressures of deuterates and the corresponding hydrates was measured in a differential tensimeter. This procedure favors cancellation of systematic error. The salts named were chosen for study be(1) The material of this article forms a portion of a thesis submitted by P. T. Miles in partial fulfilment of the requirements for the Ph.D. degree a t Princeton University. (2) PerpCrot and Schacherl, J . phys. radium, [?I.6 , 439 (1935). (3) Partington and Stratton, Nature, 189, 1076 (1936). (4) Bell, J . Chcn. Soc., 469 (1937). (6) schreherl a d Behounek, N o t w r , U8,406 (19S6). (6) M U a and M.nda. Tms Jomau., 8% 2891 (1997).
cause dissociation pressures of their hydrates are well establi~hed.~-'~ The manometric liquid employed was outgassed butyl phthalate, the suitability of which was tested by experiments which showed that significant solution and diffusion of water vapor through the manometer, or significant deuterium-hydrogen interchange, were both absent. Known weights of the salts, rendered anhydrous as described elsewhere,' were completely rehydrated within the apparatus by application of water in the vapor phase, at pressures falling short of those necessary t o yield saturated solutions. Salts thus hydrated were then eflloresced t o a composition intermediate between those of the hydrates (or deuterates) whose equilibrium pressure was desired,'S the amount of water removed for this purpose being measured in calibrated capillary tubes, later sealed off the apparatus. In a check experiment to measure the (7) Bolte, 2. physik. Chem., 86, 617 (1901). (8)Menzies, THIS JOWBNAL., 42, 1961 (1920). (9) B a x t u and Lansing, ibid., 42, 419 (1920). (10) Carpentex and Jette, ibid., 46, 578 (1923). (11) Schumb, ibid., 46, 341 (1923). (12) Menzies and Hitchcock, J . Phys. Chrm., 86, 1660 (1931). (13) Logan, ibid., 86,1035 (1932). (14) Collin!, and Mend-, ibid.. 40, 879 (19%). (16) The quwtion of the tetxahydrate of cupric sulfate portulatod by T.I. Taylor mnd othar will be discwed ebewhatr.
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FRANCIS T. MILESAND ALANW. C. MENZIES
88
Vol. 00
SrClz.6HaO SrCIr2Hr0 4- 4H10as C and to the corresponding deuterate system as D. The Heats Concerned.-The heats of dissociation for the above reactions in systems B and D as evaluated from equations (1) and (2) above and from the equations of Collins and MenziesI4 for systems A and C are: A, 13.22; B, 13.48; C, 1243 and D, 13.04 kcal. for one mole of water, as obtained by applying the factor 4.575 to the coefficient of 1/T in the equations. If we deduct from these values the latent heats of vaporization17 Results a t 37.5" of the appropriate varieties of water, 10.36 In obtaining the dissociation pressures of the and 10.67 cal., respectively, we arrive at the heats deuterates by difference, we took those of the hy- of hydration of the lower hydrates by one mole of drates from the work of Collins and Menzies." liquid water. These are A, 2.86; B, 2.81; C, 2.50 Their value for the pressure from strontium chlo- and D, 2.37, the heat of this reaction being smaller ride hexahydrate dissociating a t 25' was, inci- for both deuterates than for the analogous hydrates. In agreement with this relative order of magnidentally, confirmed by an absolute measurement of pressure after cooling the deuterate side of a tude for the cases A and B, Partington and Strattensimeter to 75" and sealing it off. Since the ton3 report the heat of hydration per mole of temperature pressure relation for dissociating water for A as 2.780 and for B, 2.218 cal. On the salt hydrates can be expressed, over the range contrary, Bell4 finds the heat of dissociation it1 here studied, by a linear equation in log p and 1/T system B 9% higher than that in A4. He states within the error of experiment, observations no heat values for system A but, using our values were made at only two temperatures, 25 and for this, we arrive a t a heat of hydration for B of 50'. In Table I the pressures for intermediate 3.74 cal., to compare with 2.86 cal. for A, showing a higher heat value in the case of the deuterate. temperatures are computed by the equations The $4 data of Schacherl and Behounek5 lead 3)DzO; log p = 10.7427 - 2946.82/T CuS01-(5 to heat values similar to Bell's for system B. (1) SrCl$.(6)r 2)DaO; log p = 10 4662 - 2850.31/T Discussion.-Table I exhibits sufficiently well (2) the experimental results of this work, and makes In Table I, which shows also the results of others, clear that there is serious discordance among the and hereafter, we shall refer to the system CuS04. results of the various observers. At 25", three 5HaO 1_ CuSOr.3HpO 4- 2H20 as A, to the cor- independent investigations show excellent conresponding deuterate system as B, to the system cordance, with a pressure for system B near 6.6 mm. while our value is 7.2.Is This 8% differTABLE I ence is hardly negligible in the study of theoretiDISSOCIATION PRESSURES IN THE SYSTEMSA, B, C AND cal signiiicance, and therefore merits discussion. D IN MM. Besides making measurements on various deuA B B C D D T. * C . C. and M." M. and M . Others C.and &%.I4 M. and M. Bell' terate systems, Bell worked also with two hydrate 20 5.3 4.9 4.4 5.8 5.5 systems, one of which, dissociating strontium chlo2.5 7.8 7.2 5. 2 8 3 8.0 6.9 ride hexahydrate, has fortunately been studied, 6 . 65S3 at 25O, by some of the most reliable especially 6.5 workers in this field, to whom Bell does not refer. 6.6 30 11.3 10 5 9.285 11 9 11.5 1 0 . 7 This makes possible a comparison of the results 9.3 of others at 25' with Bell's pressure of 7.3 mm. 9.9 Bolte' obtained 8.4, Baxter and Lansingg 8.37, 35 16.1 15.1 13.7 16.9 16.4 14.9 Schumbll 8.52, and Collins and MenziesL48.35. 22.1 21.6 39 21.3 20.0 19.4 FroweinlYreported 7.5, but his results are now 40 22.8 2 1 . 4 21.0 2 3 . 6 2 3 . 1 20 4 absolute dissociation pressure over penta- and trihydrates of cupric sulfate at 25", a constant reading was reached in three hours which was a trifle below, rather than above, the accepted value of 7.8 mm. Details of procedure and precautions to be observed have been discussed elsewhere.**"*'' Temperatures were maintained constant t o *O.O3O, and were compared within 0.02" with the scale of the Reichsanstalt. Extrapolation of pressures observed using 98.1% Dz0 to values for 100% DsO was accomplished by the method of Lewis and Cornish.16 Our pressure values may well be inaccurate even beyond *0.1 mrn., owing to contingencies other than mere error of 0bservation.1~
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50 60
44.0 81.6
41.9 78.7
42.1 80.9
44.8 81.8
44.1 81.2
(16) Lewis and Cornish, TEISJOURNAL. 66,2617 (1933).
( 1 7 ) Miles and Menzies, ibid., 68, 1067 (1936). (18) First reported by Miles, Shearman and Menzies as being 8% below the pressure for system A [Nature, lS8, 121 (1936)l. 1 ' 9 ) Frowein, Z . physik Chrm., 1, 5 , 303 (1887).
Jan., 1938
DISSOCIATION PRESSURES OF SOME DEUTERATES
89
known8 to be unreliable. Htittig and SlonimZ0re- independent method, a new pair of values of t and ported 6.35, but we understand, from a very p for system B. courteous conversation, that they prefer our The measurements of Carpenter and JettelO and value to their own in this case. It therefore of Collins and Menzies14 agree in showing that, seems possible that some systematic error has for system A, the log p against 1/T line is sensibly given Bell a value 8.14 - 7.3 = 1.1 mm. low, straight over the temperature range 25 to go', and that the same error may vitiate his other -a particular case of a general truth. If we may results. The authors2 of the value 5 mm. a t assume that the same is true for system B, we 25' for system B warn us that this was merely may use the pressure data a t 50 and 96.2' to esa preliminary measurement which was later tablish a similar straight line for this system. rectified. This yields a pressure of 7.4 mm. a t 25', thus disIt has been pointed out el~ewherel*,~~ that the favoring the lower pressures reported by other measurements a t higher temperatures on such workers. systems are more likely to be accurate, and agreeRelative Pressures.-In comparing in graphiment to 0.5% with our value is found a t 50' in cal form the results of the various observers for system B in work from another laboratory.5 We system A, clarity is gained if one employs log p and may, therefore, take the pressure for this system a t 50' as established a t 20 41.9 mm. We can avoid lower tem-
with good concordance by Speranskitl 7.45. 6,. and by Collins and Menzies,14is obI tained by multiplying by the factor 1.022 the pressure computed from the 340 330 320 310 300 Reciprocal temperature X pressure of pure water and the measFig. 1,-Logarithm of relative pressure against reciprocal temperature for ured solubility, using Raoult's law the systems A, B, C, and D. and assuming two ions. For saturated solutions of cupric sulfate pentadeuterate, 1/T as variables, since it is well known that, whose pressures we measured from 20 to 50°, we through moderate temperature ranges, the graphs found a similar factor, 1.03. Thus one can com- so obtained for dissociating salt hydrates are very pute as 552 mm. the pressure of the saturated nearly rectilinear. A much more open scale will, pentadeuterate solution a t 96.2', knowing the solu- however, be secured if one ekploys, instead of bility of the salt and the vapor pressure of deuter- absolute dissociation pressures, rather the ratio ium water. Error in the transition temperature of the dissociation pressure to the vapor pressure would be of small significance, because the p-t of pure water, of the appropriate variety, a t the curves of saturated solution and of system B have same temperature. Such ratios are fractional, similar slopes in the region a t and below the transi- and will here be called relative pressures. When tion temperature. In this way we secure, by an plotted against reciprocal temperature, the logarithm of relative pressure should diverge from (20) Hiittig and Slonim, 2. anorg. allgem. Chem., 18, 65 (1929). rectilinearity no more than does that of the vapor (21) Speranski, 2.physik. Chcm., 78, 101 (1912).
pressure of water when plotted in like manner. Through the temperature range 20 to 60°, a straight line with log p and l / T as variables reproduces the experimental pressure values for either variety of water with error riot in excess of 0.5%. In Table I1 are shown values for relative pressures in systems A and B, in computing which we employed the ordinary water pressures fouiid in “International Critical Tables,”22the deuterium water pressures of Miles and Menzies,lG the pressures for system A given by Collins and M e n ~ i e s ‘and ~ the various pressures for system B as reported in Table I. ‘r.4EI.E
11
RELATIVE PRESSURES IN PERCENTAGE FOR SYSTEMS A AND R COMPUTED FROM DATAOF VARIOUS OBSERVERS I’, “ C
20 25 30 35 39 40 50
60
A B B C and M. M.and M.S and R . 30.5 32.1 29.2 32.9 35.0 32.1 35.5 3i.7 35 6 38.3 40.6 40.6 43.0
43.2 47.6 54.6
43.6 50.3 57 7
B
B
Bell
P.and 5
31.6 33.4 39.9
32.3 33.4
41.7
42 ’i 50.5 59.3
These values are graphed in the manner de(22) “I. C. T.,” Vol. 111, 1926,p. 212.
[CONTRIBUTION FROM
THE
scribed in Fig. 1, where we have added also our lines for the relative pressures in systems C and D. If the modest extrapolation, which we have shown in Fig. 1, from the values of Partington and Stratton be allowable by reason of their utilization of four significant figures in expressing pressure, we can state that the relative pressure lines for system B as derived from the experimental data of all the otlm clhsrrvers intersect the relative pressure line for system A as graphed ill every case, contrary to our finding in this regard. It will be seen that, although it is true that the relative pressure lines of the other workers pass close to the same point at 2 5 O , this is by no means the case a t other temperatures.
Summarv Dissociation pressures in the systems consisting of deuterium water vapor and cupric sulfate penta- and trideuterate and also strontium chloride hexa- and dideuterate are reported in the temperature range 2.3 to SOo, and compared with analogous hydrate pressures. The related heat values are deduced and compared with those of others, and attention is called to certain discrepancies of our values from those obtained in three investigations by other workers. PRINCETON,
N. J.
RECEIVEDNOVEMBER 15, 1937
GEORGE HERBERT JONES LABORATORIES OF THE UNIVERSITY OF CHICAGO]
The Amine Catalysis of the Dealdolization of Diacetone Alcohol BY F. H. WESTFIEIMER AND HERZL COHEN
In their study of the dealdolization of diacetone alcohol catalyzed by amines, John Miller and Kilpatrick’ observed that, at constant buffer ratio, the rate of the reaction increased with increasFrom this, they condensation is an example using these words in the .2 Previously, however, Frenchahad found that, in phenol-phenolate ion by the hydroxyl ion buffers, the rate is eontro concentration alone. T esent work was undertaken to clear up this discrepancy. se, the rate of the deddolization
chloride, by dimethylamine and dimethylammonium chloride, by trimethylamine and trimethylammonium chloride, and by triethylamine and triethylammonium chloride, all at 18’ and a t a constant ionic strength. The work of John Miller and Kilpatrick was verified and amplified with the first two amines. Molecular trimethylamine and molecular triethylamine, however, were found to be completely without effect on the rate of the reaction. The experimental results, together with the conclusions to be derived therefrom are presented in this paper. Experimental Apparatus and Edetkod.-The rate of the reaction is easily and accurately measured dilato~netrically.~The instrument employed w a s similar to that of Br6nsted and (4)
Koelichen, 2.phyn’k. Ckm.,m,129 (1800).
