2696
H. B. HETZER,R. A. ROBISSOS,. ~ N DROGERG. BATES
that S of cold-worked metals is proportional to the concentration of defects, the activation energy for the motion of dislocations (E,) could be determined for a specimen of 90% compression. From the curve shown in Fig. 2, Emis estimated to be 1.4 e.v. in the region between 140 and 200°, which agrees with the values obtained by Phillips'z (1 2 e.v. a t 200"). The recovery curve for cold-worked nickel (80% compression) has two stages, one between 200 and 300" and the other above 400" as shown in Fig. 4A. According to Clarebrough and his co-w0rkers,~-5 these are attributed to the disappearance of vacancies ( Tv) and dislocations (TD), respectively. Tammann reports that the changes in S occur only in the region from 150 to 350". 11. Lattice Defects as Active Centers in Heterogeneous Catalysts Uhara and his co-worker~~.'~ found that the activities of twisted copper for the decomposition of diazonium salt and dehydrogenation of ethyl alcohol considerably decrease by annealing at about 350°, as shown in Fig. 3 . This temperature range coincides with the recovery temperature of S (Fig. 1) for a slightly cold-worked specimen (the compression is usually more severe working than twisting as shown by the comparison of induced catalytic activity). I n 1933, it was found by Eckell14that the catalytic activity of nickel for the hydrogenation of ethylene was enhanced by a factor of 600 to 1000 after rolling. Since the activity is eliminated by annealing in the region from 200 to 300°, as shown in Fig. 4B, we may conclude that the active center
-
(12) V . A. Phillips, J . Inst. Metals, 81, 185 (1952). (13) I. Uhara, S. Yanagirnoto, K. Tani, G. Adachi, and S.Teratani, J . Phys Chem, 66, 2691 (1962). (14) J. Eckell, 2. Eiektroeheni., 39, 433 (1933).
Vol. 66
is some kind of point defect a t the surface, contrary to Cratty and Granato's postulate15 ascribing the activity to dislocations. Uhara, et. aL,8 found that the catalytic activities of cold-worked nickel for para-ortho (p-0) Hz conversion, dehydrogenation of ethyl alcohol, and hydrogenation of cinnamic acid decrease suddenly in the two temperature ranges Tv and T D , perfectly in parallel with the change in S as shown in Fig. 4 C 4 E , when the measurements were carried out with the same cold-worked specimens. Consequently, active centers in these catalysts may be both point defects and the termination of dislocations a t the surface. iiccording to Cremer and Kerber,16 the activity of nickel foil for p-o Ht conversion was decreased by raising the temperature of annealing as shown in Fig. 4F. Since the foil was prepared by severe deformation, it is felt that T Dwas shifted to lower temperatures, eventually overlapping Tv. Similarly, it is almost impossible to find the distinction between TV and T D for active centers in ordinary catalysts, which are prepared by chemical procedures and contain a large number of defects, as shown by X-ray studies. The measurement of S of cold-worked metals can be carried out readily by a simple technique with a small quantity of the sample, and hence it offers a very convenient method for identifying lattice defects with active centers in metallic catalysts. Studies on platinum catalysts will be reported in forthcoming papers. Acknowledgment.-The author wishes to express sincere thanks to professor Uhara for his kind guidance and to Dr. Saika for his advice. (15) L. E. Crattyand A. V. Granato, J. Chem. Phys., 26,96 (1957). (16) E. Cremer and R. Kerber, Advan. Catalysis,7 , 82 (1955).
DISSOCIATION CONSTANT OF t-BUTYLAJIMONIUM ION AND RELATED THERMODYNAMIC QUAXTITIES FROM 5 TO 35' BY HANNAH B. HETZER, R. A. ROBINSON, .4ND ROGER G. BATES Solution Chemistry Section, National Bureau of Standards, Washington 26, D. C. Received June 26, 1963
The acidic dissociation constant of t-butylammonium ion has been determined from 5 to 35" by e.m.f. measurements of hydrogen-silver bromide cells without liquid junction. At 25", -log Kt,h = 10.685, and the temperature coefficient of the dissociation constant gives the values AH0 = 60,070 j. mole-' and ASo = -3.1 j. mole-' deg.-'. These thermodynamic constants are compared with the corresponding values for the acidic dissociation of the protonated forms of the aminoalcohols related structurally to t-butylamine.
Introduction The relationship between basic strength and molecular structure has long been of interest.' Recent studies of the basic dissociation of the aminoalcohols have shed further light on the influence of structural factors in this class of com-
pounds. To this end, the dissociation constants for the three ethanol-ammonium ions were deterIn addition, the three structurally related substituted ammonium ions, namely, 2ammonium- 2 - (hydroxymethyl) - 1,3- pr~panediol,~ 2-ammonium-2-methyl-1,3-propanediol,6~7 and 2-
(2) Monoethanolamrnonium: R. G. Bates and G. D. Pinching, (1) See, for example, D. H. Everett and W. F. K. Wynne-Jones, J. Res. Natl. Bur. Std., 46, 349 (1951). Proc. Roy. Soe. (London), A177, 499 (1941); A. G. Evans and S. D. (3) Diethanolammonium: V. E. Bower, R. 9.Robinson, and R. G. Hamann, Trans. Faraday Soc., 47, 34 (1951); D. H. Everett and B. R. Bates, ibid., 66A, 71 (1962). W . Pinsent, Proc. Roy. SOC.(London), U 1 6 , 416 (1952); R. P. Bell, (4) Triethanolarnrnonium: R,G. Bates and G. F. Allen, ibid., M A , "The Proton i n Chemistry," CornelJ University Press, Ithaea, N. Y., 343 (1960). 4959 chapter 5.
DIBSOCIATIOS COSST.AST OF ~-BUTYLAMMONIUM Iox
Dec., 1962
2697
TABLE I ELECTROMOTIVE FORCE OF THE CELL: P t ; H?(G.,1 ATM.)(HsC)&NHsBr(m1),(H3C)3CNH2(m2), AgBr; Ag,
FROM
5
TO
35"
(IN V.) ml
ma
50
100
0.09065 ,08043 ,07316 ,07081 ,06422 ,05890 ,05350 ,05057 ,04724 .04256 .04045 .03592 .03374 ,02421 ,01293
0.04588 ,03945 ,03702 ,03536 .O x 5 0 ,02942 ,02708 ,02516 .02350 ,02126 ,01983 ,01818 ,01655 ,01225 ,006341
0.76420 ,76562 ,76843 ,76900 ,77016 , 77243 ,77449 ,77543 ,77678 ,77887 ,77927 ,78231 33277 .79022 ,80205
0.76357 ,76502 ,767117 ,76833 ,76959 ,77196 . 77404 r--
.iio04 ,77641 ,77846 ,77893 ,78207 ,78252 .78999 .80203
200
25a
0.76173 ,76331 ,76620 ,76665 ,76793 , 77039 , 77248 ,77351
0.76058 ,76226 ,76504 ,76555 ,76686 , 76935 , 77148 ,77255 ,77405 ,77625 ,77681 .77996 ,78049 ,78838 ,80096
15'
0.76271 ,76427 ,76711 ,767X ,76884 . 77126
,77335 ,77434 rF--' ,iioi3 ,77782
,77494 ,77708
.7i834 ,78149 .78199 ,78959 ,80181
ammonium-2-methyl-1-propanol7 have been studied. Data now have been obtained for the fourth member of this series, namely, the hydroxyfree parent compound, t-butylamine.
,77763 .78081 ,78132 .78901 ,80149
Pt; Hz(g., 1 atm.), (H3C)3CXH3Br(rnl), (H3C)3CXH2(m2), AgBr; Ag where m is molality. The preparation of the standard solution of hydrobromic acid and of the silver-silver bromide electrodes has been described p r e v i ~ u s l y . ~ A commercial preparation of t-butylamine waa distilled ( a t approximately 25" and 400 mm.) through a spinningband distillation column of the Piros-Glover type.I0 The column had a rectifying section about 60 cm. in length and was rated at about 80 theoretical plates. Several fractions (about 17 in number) were analyzed by vapor-liquid chromatography, and the best were retained for use. The commercial amine gave a band believed to indicate the resence of a higher-boiling im urity, but the material u s e l f o r the e.m.f. measurements &owed no peak on the chromatogram except that for t-butylamine. In order t o correct the observed e.m.f. values t o a pressure of 1 atm. of dry hydrogen, the partial pressure of the amine from 0.1 M and 0.05 M aqueous solutions was measured a t 25,40,45, and 50" by a gas-transpiration method." Henry's law appeared to be valid over this limited range of concentrations. The constant, k = p/m, for partial pressures in mm., was found t o be approximately 20 a t 25", 64 a t 40°, 82 a t Go, and 123 a t 50". Values of k a t 30 and 35" were read
350
0.75783 ,75960 ,76243 .7629:3 , 7643:3 ,76690 ,76900 ,77014 .77171 ,77396 ,77462 ,77784
,77837 .78654 ,79925
TABLE I1
VALUES OF --LOG
Kb
K b h AND -LOG
Temp., "C.
-log Kbh
5 10 15 20 25 30 35
11.439 11.240 11.048 10.862 10.685 10.511 10,341
Experimental Method and Procedure .-The electromotive force method used was essentially that previously employed in this Laboratory for studies of other volatile bases.8 Because of the considerable solubility of silver chloride in solutions containing t-butylamine, however, the silver-silver bromide electrode was substituted for the silver-silver chloride electrode. The cell used, therefore, is represented as
30'
0.75927 ,76101 ,76381 ,764335 ,76587 ,76820 , 77033 ,77143 ,77295 ,77518 .77586 .77896 .77950 .78753 .80042
FROM
5 TO 35" -log Kb
cl
3,295 3.295 3.298
0,002 ,001 ,001 ,001 .001
3.305 3.311 3.322 3.339
.001 .002
TABLE I11 THERMODYNAMIC QUANTITIES FOR THE ACIDICDISSOCIATION O F f-BUTYLAMMONIUM I O N (BH') FROM 5 TO 35" t.
oc. 5 10 15 20 25 30 35
AGO, j . mole-'
AHQ, j. mole-'
60,910 60,930 60,950 60,970 60,980 61,000 61,010
59,780 59,850 59,920 59,990 60,070 60,140 60,220
1.
ASO, deg.-l mole-'
-4.1 -3.8 -3.6 -3.3 -3.1 -2.8 -2.6
from a plot of log k against 1/T, where "5 was in "K. It was found necessary t o take the vapor pressure of the amine into account only a t 35" and for buffer solutions in which mz was 0.025 or greater. For other temperatures and concentrations, the pressure corrections calculated with and without the partial pressure of the amine agreed t o within 0.02 mv., and the contribution of the amine to the total pressure was neglected accordingly. To reduce losses by volatilization, the ratio ml/mz was kept a t about 2 in all of the cell solutions. Values of Kr, the stability constant for the diammine complex formed by silver ion and t-butylamine, were calculated from measurements of the solubility of silver chloride (5) R. G. Bates and H. B. Hetzer, J . P h y s . Chem., 65, 667 (1961). in 0.11 M and 0.07 M solutions of the amine. The results for log &were as follows: 8.78 ( O " ) , 7.88 (25O), 7.12 (50"). (6) H. B. Hetzer and R. G . Bates, ibid., 66, 308 (1962). The solubility product of silver chloride being 2,500 times (7) B. A. Timini, thesis, University of Bristol, 1960; D.H.Everett, greater than that of silver bromide ( a t 25O), it can be shown private communication. that the solubility of silver bromide in the buffer solutions (8) See, for example, R. G. Rates and G . D. Pinching, J . Res. IVatZ. would be too small to require corrections to the stoichioBur. Std.. 4'2, 419 (1949). metric bromide concentration (ml). Corrections to ml and (9) H.€3. Hetzer, R. A. Robinson, and R . G. Bates, J . Phys. Chem., m 2 for the dissociation of the amine were necessary, how66, 1423 (1962). This paper also gives the values of the standard ever, and were applied in the usual way. The molality of e.m.f. of the cell which was used in deriving the dissociation constant free amine in the stock solutions was determined by weight from the e.m.f. data. titration t o the calculated equivalence point (5.9 in 0.05 If (10) The authors are indebted to Dr. R. T. Leslie and hlr. E. C. solution), detected by means of a glass electrode. Kuehner for the distillation and for the vapor-liauid chromatographie Because of the high volatility of the base, the e.m.f. analysis of the amine. decreased slowly with time. After the completion of a series (11) See ref. 8,
H. B. HETZER, R. A. ROBIWSON, AND ROGER G. BATES
2608
COMPARISON OF T H E
TABLE IV THERMODYNAMIC QUANTITIES A T 25' F O R T H E DISSOCIATION O F t-BUTYLAMMONIUM STRUCTURALLY RELATED AMMONIUM ALCOHOLS M.P. of
Base
Ref.
( CH20H)&. N H ~
-
AGO,
I O N AND
ASQ,
THREE
ACpO,
-log Kbh
j . mole-'
AHQ, j. mole-'
..
8.075 8.801 8,790
46,075 50;238 50,176
47,600 49; 860 49,720
5.1 -1.3 -1.5
- 64 -45 - 46
30-3 1
9.691
55,318
53, '330
-4.6
-25
10.685
60,980
60,070
-3.1
15
amine, OC.
1'71 111
(CH20H)&.~~? 6
I
Vol. 66
j. deg.-' mole-' j .
deg.-l mole--
CH, CH; CH20H--L-~H2
I
7
CH 3 ( C&)&
-67.5
".2
of Pearson and Williams'2; 10.937 at 18" with 10.83 from the measurements of Girault-Vexlearschila; 10.685 a t 25" with 10.53 from the conductance data of Bredig.I4 The thermodynamic quantities for the dissociation of t-butylammonium ion in the standard state were derived from the constants of eq. 1and are listed in Table 111. Only an approximate value of AC,O can be calculated from data over a short temperature range, but eq. 1 gives ACPo = 15 j. deg.-l mole-' at 25'. The following values Results pertain to the basic dissociation of t-butylamine The corrected values of the e.m.f. of the cell are at 25" given in Table I. The calculation of the acidic Process: (H3C)&?;H2 HzO dissociation constant, Kbh, of the t-butylammonium (HSC)3CNH3+ OHion was made in the manner already described.6s6 A value of 0 for the ion-size parameter (a*) gave AGO = 18,910 j. mole-' straight-line plots of K b h ' as a function of ionic A H o = -3-150 j. mole-' strength at each temperature. The intercepts were ASa = -73.0 j . mole-l deg.-l determined by the method of least squares. The values of -log K b h and of -log Kb (the negative The thermodynamic quantities for the dissocialogarithm of the basic dissociation constant of tion of t-butylammonium ion are compared in t-butylamine) are collected in Table 11; is the Table IV with the corresponding quantities for the standard deviation of the intercept (-log K b h ) protonated forms of the three aminoalcohols reat each temperature. lated structurally to t-butylamine. A small unThe following equations represent the "observed" explained irregularity in the trend of ASo is apdata at the seven temperatures with a mean devia- parent, but the change of entropy is generally in tion of *0.001 unit qualitative agreement with expectation, insofar as the effect of hydroxyl groups on the entropy of 3021.63 solvation is concerned. The melting points show -log K b h = ___ 0.9376 - 0.001303T T that the successive introduction of hydroxyl into the methyl groups of the parent compound produces (1) a considerable lowering of the volatility of the amine 1477.29 as well. -1OgKb = - 7.2023 -k 0.018647T (2) of measurements, the e.m.f. of test cells without the triple saturator for the hydrogen gas was found to be about 1 mv. lower than a t the beginning of the series, some 52 hr. earlier. The saturators were used accordingly for all of the measurements from which the dissociation constant was derived, but decreases of e.m.f. rangingfrom0.25 to0.5mv.,roughlyproportional t o elapsed time, still were observed. The e.m.f. values at temperatures ot,her than 25" therefore were corrected t o "zero time" in order to allow for loss of the amine through volatilization.11 The lugest correction is a t 35", where it amounts to 0.008 in (-log K b h ) ; a correction of this small magnitude can be made with confidence.
+
+
+
~
T
The value of -log K b h = 11.644 at 0" (calculated by eq. 1) can be compared with 11.64 from the work
(12) R. G. Pearson and E'. V. Williams, J . A m . Chem. SOC..76, 258 (1954). (13) G. Girault-Vexlearschi, Bull. soc. chim. Francs, 577 (1956). (14) G. Bredig, 2. physik. Chem., 13, 289 (1894).
