Anal. Calcd. for C12HCLo02Cs: C, 21.69; H, 0.15; C1, 53.35. Found: C, 21.88; H , 0.17; C1, 53.16. Cesium Hydrogen Di-(3-nitrophenoxide), CsH( NO&aHdO)z.This compound was obtained as large transparent ruby rhombohedra' using the above procedure, or by using ether in place of benzene. Anal. Calcd. for C12H&~O&s: c , 35.14; H, 2.21; N,6.83. Found: C, 38.32; H , 2.48; N , 6.97. Sodium Hydrogen Di-(4-nitrophenoxide) Dihydrate, NaH (NOZC~H~O)Z.~HZO.--A solution containing 10 mmoles of 4nitrophenol in 3 ml. of methanol was treated with 5 ml. of 1 12' YaOH. The resulting orange needles6 were twice recrystallized from water. Anal. Calcd. for C1&N208Na: C, 42.86; H, 3.90; N, 8.33; E a , 6.84. Found: C, 42.93; H , 4.25; N, 8.38; Na, 6.41. Potassium hydrogen di(4-nitrophtnoxide) dihydrate, KH( K O Z C ~ H ~ O ) ~ . ~was H Zprepared O, as above using KOH instead of NaOH . Anal. Colcd.: neut. equiv., 352.3. Found: neut. equiv., 350.8. Potassium Hydrogen Di-(4-cyanophenoxide), KH( CNC6H40)2. -To a solution of 2 mmoles of 4-cyanophenol (Aldrich Chemical Co.; three times crystallized from Chloroform-carbon tetrachloride and once from water; m.p. 112" (liteao m.p. 113")) in 12 ml. of anhydrous ether waa added 1 ml. of 1 N methanolic potassium hydroxide. White crystals of acid phenoxide precipitated. This substance was a monometbanolate as found by analyais; drying of the compound for 10 hr. a t 75" and loF4 mm. gave unsolvated acid phenoxide. Anal. Calcd. for C14H9N~02K: C, 59.07; H , 3.43; X, 10.60; K, 14.79. Found: C, 59.29; H , 3.57; K(Kjeldahl), 10.29; K, 14.62. Potassium Hydrogen Pyrocatechoxide Hemihydrate, KO( O H ) C ~ H ~ . ~ / Q H ~ O . -compound T~~S was prepared from pyrocatechol (Eastman Kodak Vhite Label) and a 157, solution of potassium hydroxide in ethanol using Weinland's method .lo Anal. Calcd.: neut. equiv., 157.2. Found: neut. equiv., 158.0. Potassium trihydrogen pyrocatechoxide hemihydrate, KO(OH)C6H4.C6H4( OH)2J/2Hz0,was prepared similarly as above using 1 Ai KOH.'O Anal. Calcd.: neut. equiv., 267.3. Found: neut. equiv., 266.5. Potassium Trihydrogen Guaiacoxide, KOC6H40CH8-3a 10% ethanolic solution of guaiacol HDCeH40CH3-To (Matheson Coleman & Bell Co.), 1 N KOH was slowly added. White crystals deposited and were recrystallized from water. Anal. Calcd.: neut. equiv., 534.5. Found: neut. equiv., 530.1,
Sodium acid salt of 2,2'-dihydroxybiphenyl, NaOC6H4G&OH, was prepared by careful half-neutraliaation of phenol solution in ethanol with 1 AT NaOH. White crystals deposited from the solutian which was concentrated by evaporation. Aldrich Chemical Co. 2,2'-dihydroxybiphenyl was crystallized 5 times from toluene, m.p. 108-109.5° (lit.31m.p. loso), neut. equiv.: calcd. 186.2; found 186.6. Sodium acid salt of 2,Z.'-methylenebis-(4-chlorophenol), KaOCaHs(Cl)-CHz-(Cl)COHaOH, was prepared as above. 2,2'Methylenehis-( 4-chlorophenol) (British Drug House) was crystallized four times from ethylene chloride, m.p. 187-169" after sintering at 162'; neut. equiv. calcd, 268; found 270. Adduct of 4-Nitrophenol with Potassium Acetate, NOzCeH40H. MOOCCHs.-A splution containing 20 mmoles of 4-nitrophenol in 50 ml. of ethanol was treated with 20 mmoles of potassium acetate dissolved in water. Yellow crystals (m.p. 327") deposited after a few hours and were recrystallized from ethanol. Anal. Calcd.: neut. equiv., 237.2. Found: neut. equiv., 240.1. Adduct of 4-Nitrophenol with Trimethylamine Oxide, NO2C6HIOH.NO(CHa)8.-This adduct was prepared by adding a solution of 4-nitrophenol in ethanol to a solution of trimethylamine oxide in equimolar ratio, Yellow crystals (m.p. 115") deposited and were recrystallized from ethanol. Tripethylamine oxide dihydrate was synthesized according to reference 32. Adduct of Pyrocatechol with Potassium Oxalate, 2CBH4( OH)p K2C204.-This adduct was prepared by adding an aqueous solution of pyrocatechol (Eastman Kodak White Label) to potassium oxalate solution (molar ratio 2: 1). White crystals deposited and were recrystallized from water. Deuterated Compounds.-Substancw were deuterated on the vacuum line by exchange with methanol-d or deuterium oxide. Dimethyl sulfoxide (Light and Co.) was distilled under reduced pressure from molecular sieves (type 4-A, Linde Air Products). Infrared Spectra.-Substances were prepared for spectroscopical work as mulls in Nujol and hexachlorohutadiene. Some adducts were investigated a t liquid nitrogen temperature as mulls in Nujol; a law temperature cell of a conventional design was used for this work. About 207, solutions of adducts and of phenols in dimethyl sulfoxide (DMSO) were studied as thin films hetween NaCl plates using the compensation method. The spectra were recorded on Perkin-Elmer Model 21 and Beckman IR-4 spectrophotometers equipped with sodium chloride prisms.
Acknowledgment.-D, H. and ,4.N. are grateful to the Boris Kidri6 Fund for a substantial contribution toward the expenses of experimental work. (31) W. R. Spencer and F. R. Duke, Anal. Chem.. 26, 919 (1954). (32) J. Meisenheimer and K. Bratring, Ann., 597, 286 (1913).
(30) 0. L. Brady and F. P. Dunn, J. Chem. Soc., 871 (1914).
DISSOCIATION COXSTAST OF PYRROLIDlNlURl 10s ASD RELATED THERMODYNAMIC QUAXTITIES FROM 0 TO 50' B Y HAKNAH B. HETZER, ROGERG.
BATES, AND
R. A. ROBISSON
Solutzon Chemistry Section, National Bureau o j Standards, Washington, L). C. Received December 6 , 196.8
The thermodynamic dissociation congtant of pyrrolidinium ion (BH') a t 11 temperatures from 0 to 50" has been determined from e.m.f. measurements of hydrogen-silver bromide cells without liquid junction. The disHaO+ is given as a function of T (OK.) by the sociation constant ( K b h ) for the process BH+ H20 + B 5.2942 - 0.00592321. At 25", -log Kbh is 11.305, AHo is 54,470 j. mole-', equation -log Kbh = 2318.85/T ASo is -33.7 j. deg.-l mole-', and ACPois 68 j . deg.-' mole-I.
+
+
Introduction In an earlier contribution from this Laboratory,+ the acidic dissociation constant of piperidiiiium ion from 0 to 50' was reported. This paper describes a continuation of work on the dissociation constants of nitrogenous bases by extending measurements to the next lower member of the series, pyrrolidine (tetra(1) R . G. Bates and T. E. Bower, J. Res. A'atl. Bur. Std., 67, 153 (195G).
+
methyleneimine (CHz)4NH), Apart from the interest in a comparison of the thermodynamic properties of these two closely related bases, pyrrolidine is important in being the parent of the amino a i d s proline and hydroxyproline, whose dissociation constants have already been measured.2 Moroever, the ring structure of four ( 2 ) P. K. Smlth, A T. Gmham, and E. R. B. Smith J . Bzol. Chem , 144, 735 (1942).
DISSOCIATION COXSTANT
May, 1963
1325
O F PYRROLIDINIUM I O N
TABLE I
ELE:CTROMOTIVE FORCE OF 77h1
0.09946 ,09520 ,08348 ,08296 ,07904 ,07409 ,07239 .06928 ,06741 ,06482 ,06169 ,05568 .05062 ,04797 .04395 ,03865 ,03080 ,02918 .02624 ,01922
m2
00
THE 50
CELLPt; Hg(g, 1 ATM.), (CH-JrNHzBr(ml), (CH&NH(mz), AgBr; Ag 100
15'
250
30'
350
FROM
40'
0 TO 50' 450
(IN
v.) 50'
0.05042 0.79210 0.79330 0.79423 0.79495 0.79542 0.79580 0.79584 0.79573 0.79555 0.79509 0.79454 ,79581 ,04954 .79333 ,79455 ,79556 ,79633 ,79688 ,79721 ,79738 ,79731 .79708 .79658 ... ... ... ~ . . ... ,79626 ,79644 ,79638 ,79616 ,79573 ,7951'7 ,03799 ... ... ... ,04024 ,79411 .79532 ,79635 . '9711 ,79763 ,79802 ... ... .04006 ,79587 ,79717 ,79822 .79898 ,79958 .79990 .80002 ,80001 ,79983 ,79961 ,79913 ,797413 ,03372 ... ... ... ".. ... ,79843 ,79863 ,79862 ,79842 .79809 ... ... ... ,03511 ,79644 ,79771 ,79875 ,79956 ,80015 ,80057 ... ... ,80301 ,80300 ,03605 ,79883 ,80013 ,80123 ,80208 ,80271 .80313 ,80338 ,80344 .80335 ,79996 .79943 .03068 ... ... ... ... ... ,80022 ,80048 .80050 .80034 ... ... ... ... ... ,03144 ,79824 .79947 ,80062 ,80151 ,80214 ,80254 ,80478 ,80434 ,03127 ,80035 ,80169 ,80276 ,80359 ,80419 ,80468 ,80495 ,80505 ,80509 ,02897 ,80269 ,80403 ,80522 ,80620 ,80690 ,80729 ,80772 ,80780 ,80786 ,80763 ,80720 805611 ,80623 ,80600 ,02303 .. , ... ... ... ... ,80573 ,80609 .80625 ... ... ... ,02326 ,80365 ,80507 ,80624 ,80720 ,80793 ,80847 ... ... ,81254 .81226 .02287 ,80684 ,80831 ,80954 .81057 ,81133 ,81187 ,81229 .81251 .81264 ,01959 ... .81032 ,81156 ,81253 ,81327 ,81385 ,81426 ,81458 ,81468 ,81458 ,81433 ,01561 ,81273 ,81430 ,81560 ,81669 ,81756 ,81814 ,81857 ,81894 ,81911 ,81905 ,81891 ,01328 ... ... ... ... .81632 ,81681 .81711 .81727 ,81722 ,81698 ... ... ... ,01273 ,81443 ,81599 ,81732 ,81843 ,81933 .82004 ... ... .01000 ,82152 ,82306 ,82459 .82580 ,82684 ,82759 ,82830 ,82877 ,82917 ,82936 ,82939
carbon atoms and one basic nitrogen atom occurs in many natural products, such as the alkaloids hygrine and nicotine, chlorophyll, and hemoglobin. Method and Procedure The electromotive force method employed was essentially that used in previous studies of bases in this Laboratory.l*3-j As in the study of t-butylamine,s the silver-silver bromide electrode, because of its lower solubility in amine solutions, was used in place of the silver-silver chloride electrode. The cell uved may be represented as
where m is molality. The preparation of the electrodes and of the hydrobromic acid solution have been described elsewhere, and the values used for the standard potential of the cell have already been reported.6 Two separate lots of pyrrolidine, commercial grade, were distilled through a Podbielniak column with platinum Heligrid packing, under an atmosphere of nitrogen. The middle fractions were ana,lyzed chromatographically, and the fractions used were those for which the chromatogram showed only the pyrrolidine peak. Analysis of a sample from the first distillation with the mass spectrograph revealed only one component .7 Although the purified amine was protected from moisture and carbon dioxide during storage, evidence of decomposition (slight initial instability of the electrodes and lower e.m.f. values) was observed in buffer solutions prepared from purified pyrrolidine several weeks after the distillation had been carried out.5 There was no evidence, however, of decomposition of the base used to prepare the buffer solutions for the determination of the dissociation constant. Since pyrrolidine is rather volatile (b.p. 86.5" at 760 mm.), the ratio of salt to free base (ml/mz) in the buffer solutions was kept a t approximately 2 . The molality of free pyrrolidine (mz) in the stock solutions was determined by weight titration with the standard solution of hydrobromic acid to the calculated equivalence point (pH 6.1 in 0.1 M solution), the end point being determined with a glass electrode. Insofar as possible, all of the solutions containing pyrrolidine were prepared and handled under an atmosphere of nitrogen. (3) R. G. Bates a n d G. D. Pinching, J . Res. Nutl. Bur. Std., 42, 419 (1949)' N.B. Hetzer a n d R. G. Bates, J . Phys. Chem., 66, 308 (1962). ( 5 ) H. B. Hetzer, R. A. Robinson, and R. G. Bates, i b i d . , 66, 2696 (1962). (6) H. B. Hetzer, R. A. Robinson, a n d R . G . Bates, ibid.,66, 1423 (1962). (7) The authors are indebted to Dr. R. T. Leslie for the distillation, to Dr. Leslie and E. L. Weise f o r the vapor-liquid chromatographic analyses, and to E. E. Hughes for the mass spectrographic examination. (8) Decomposition of pyrrolidine has been observed by R. V. Helm, W. J. Lanum, G. L. Cook, and J. S.Ball, J . Phys. Chem., 62, 858 (195S), and R. H. Linnell, M. Aldo, and F. Raah, J . Chem. Phys., 86, 1401 (1962). (4)
200
The vapor pressure. of pyrrolidine from the most concentrated 0.1; m2 = 0.05, was measured a t BO' buffer solution, ml = by a gas transpiration method and found to be approximately 0.5 mm. The pressure corrections to one atmosphere of dry hydrogen, calculated with and without inclusion of the partial pressure of the amine, agreed to within 0.01 mv. The partial pressure of the amine was therefore neglected. Although the cells were equipped with the extra triple s a h rators used in this Laboratory for solutions with volatile components,s the e.m.f. values a t 25' measured in the middle of the series and a t the end were lower than the original values. This decrease was found to be approximately 0.1 mv. after the measurements from 0 to 20" and approximately 0.4 mv. after those from 30 to 50'. The e.m.f. a t temperatures other than 2'5" was therefore adjusted by amounts roughly proportional to time and concentration, determined from the changes in the successive readings a t 25'. The largest adjustment made (that for the solution with ml = 0.03080 a t 50') amounted to 0.008 p K unit. No data on the stability of the pyrrolidine complex with silver ion were found. Even if the stability constant were considerably higher than that for the ammonia and piperidine complexes, however, the low solubility of silver bromide would keep the concentration of complex very low. Consequently, no correction to the stoichiometric concentration of bromide was necessary. Hydrolysis corrections to the buffer ratio were made in the usual way. N
N
Results The e.m.f. data are listed in Table I. Each value represents the mean of two electrode combinations in the same cell, corrected as described in the previous section for removal of the volatile base by the bubbling hydrogen. The usual correction to a hydrogen partial pressure of 1atm. has also been made. The calculation of K b h , the thermodynamic dissociation constant of the pyrrolidinium ion (BH+) on the molal scale, was made in the manner alreadly de~cribed.~Once again, a value of 0 for the ion-size parameter gave a crtraight-line plot of the "apparent," log K'bh, namely log K ' b h , as a function of ionic strengtlh, I . The equation for the calculation of -log K ' b h w a ~ therefore
(9) R. G. Bates and H. B . Hetzer, J . Phys. C'henz.. 65, 667 11961:.
H. B. HETZER, R. G. BATES, A S D R. A. ROBISSOT
1126 12.18)
Vol. 67 TABLE I1
I
I7.4LUES O F -LOG
K b b AND -LOG
Kb FROX
0 TO 50"
t , "C.
-log K b h
ai
-108 Kb
0 5
12,165 11.984 11.807 11.636 11.469 11.305 11,147 10.994 10.845 10.698 10.556
0,001 002 ,001 ,001 ,001 ,001 ,001 ,001 ,001 .OOl .002
2.779 2,750 2.728 2.711 2.698 2.692 2,686 2.686 2.690 2.698 2.706
10
15 20 25 30 35 40 45 50
t
and -log Kh l
I
I
I
10.53 -
10.49 -
I
,
0
0 05
of -log K", as a function of ionic strength ( I = ml m o p ) at 0, 25, and 50".
+
BH+
+ E120 J_ B +
H 3 0 +
+
-log
Kbh
=
2318.85 ___ + T
5.2942 - 0.005923T
(4)
(10) R. G. Bates, G. L. Siesel, and S. F. Bcree, J . Res. T a i l . Bur. Std., 51, 205 (1943). (11) H. S. Earned and R. A. Robinson, Trans. Faradag Soc., 36, 973 (1940).
=
15,350 j. mole-1
AHo
=
2,110 j. mole-]
=
-44.4 j . deg.-l mole-1
ACPo= -270 j , deg.-' mole-1 TABLE 111 THERMODYNANIC QUAKTITIEB FOR THE ACIDICDISSOCIATIOX OF PYRROLIDISIUM ION(BH") FROM 0 TO 50" t,
oc.
0 5 10 15
20 25 30 35 40
45 50 25":
that is, the acidic dissociation of pyrrolidinium ion, whereas K b is the equilibrium constant for the rcaction
+
+ 0.023461T
AGO
AXo
(2)
B H,O BH+ OH(3) The product of these two constants is K,, the autoprotolysis constant for water, values for which were taken from the paper of Harned aiid Robinson.'l Thermodynamic Quantities.-The following equations, the constants of which were determined by the method of least squares, represent the values of -log Kbh and -log Kb with a mean deviation less than 0.001 unit
~
+
0 10
where A is the Debye-Huckel slope. The concentration of hydroxide ion was estimated from the values of p(mn~c1)and pK, in the manner suggested by Bates, Siegel, aiid Acree.lo The values of -log K ' b h a t 0, 25, and 50' are plotted as a function of ionic strength in Fig. 1. The intercepts (-log K b h ) , summarized in Table 11, were determined by the method of least squares, and the standard deviation of the intercept (ul) is given in the third column of the table. Values of -log Kb, the negative logarithm of the basic dissociation constant, are given in the last column. The constant K b h is the equilibrium constant for the reaction
2196.67 - 11.6690 T
where T is the temperature in degrees Kelvin (2'C. 273.15). By application of the usual thermodynamic equations t'o the expression given in eq. 4, the standard thermodynamic quantities for the dissociation of pyrrolidinium ion (eq. 2 ) given in Table I11 have been calculated. The corresponding quantities for the basic dissociation (eq. 3) at 25' are
1.
Fig. 1.-Plot
=
AGO, j. mole-'
AHQ, j . mole - 1
63,620 52,850 53,170 63 ,810 53,480 64 000 53,810 64,180 54,140 64,360 54,470 64 , 530 54,810 64,700 55,160 64,860 55,510 65,010 55 ,870 65,160 56,230 65,310 ACpO= 68 j . deg.-l molew1
AS0 .
i.d e s -1 mole -1
-39.4 -38.3 -37,1
-36.0 -34.9 -33.7 -32.6 -31.5 -30.3 -29.2 -28.1
Discussion Earlier determinations of the dissociation constant of pyrrolidinium ion all seem to have been based on the e.m.f. of cells with liquid junction (l.j.).'z-ls These values are summarized in Table IV. The values of AHa and ASo at 30' calculated by Searles, Tamres, Block, and QuartermanI4 from their dissociation constants at 25 and 35' are greatly a t variance with those found in this investigation. These values are 33,500 j . mole-l and -104 j . deg.-l mole-l, R. M. Hixon, J . Am. Chem. SOC.,68, 4367 (1931). (13) A. Albert, "Heterocyclic Chemistry," Athlone Press, London, 1959. (12) L. C. Crais and p. 6.
(14) S. Searles, M. Tamres, F.Block, and L. A. Quarterman, J. Am. Chem. S o c . , 78, 4917 (1986).
( 1 5 ) J. A. Broomhead, H. A. McKenrie, and D. P.Mellor, Australian J . Chem., 14, 649 (1961).
DISSOCIATION CONSTANT OF PYRROLIDIKIUM 10s
May, 1963
respectively. On the other hand, the heat of neutralization of pyrrolidine with hydrochloric acid in aqueeous solution at 25’ determined by Facconi, Paoletti, and CiampolinilBleads to a value of 51,750 j . mole-1 for AHo, in reasonably good agreement with 54,570 j . mole-1 calculated from the change of K h h with temperature. These authors have given a value of 38.9 j . deg.-l mole-I for ALYO,~’ but, combined with our values for AGO at 2 5 O , their calorimetric data lead to AX0 = -42.9 j . deg.-l mole-’. This figure i s to be compared with -33.7 j . deg.-l mole-1 listed in Table 111. COMPARIbON t,
“C.
..
OF
Ionio strength
TABLE IV VALUES FOR -LOG K b h (PYRROLIDINIUM -log Kbh
Method
.,.
11 11’ Hz, 1.j. 20 11.3 .... Glass,1.j. 25 0 11.27 HP,AgBr 0 11 305 Glass, l.j. 0 2 114 35 0 11.08 Glass, l.j. 0 10,994 H2,AgBr a If temperature is taken to be 25’.
ION)
Ref.
12 13 14 This investigation 15 14 This investigation
Pyrrolidinium ion is a slightly weaker acid than piperidinium ion, with which it is closely related structurally. The enthalpy and entropy changes 011 dissociation of the two acids are strikingly similar; the changes of heat capacity are rather large and positive for both acids. AHO,
Acxd
--log K b h
J.
Piperidiniuni Pyrrolidinium
11 123 11.305
53,390 54,470
mole-’
ASO, j . deg. -1
A Cp,‘ 1. den. -1
mole-‘
mole-’
-33 9 -33 7
88 68
The data, are for a temperature of 25’. The presence of a carboxyl group adjacent to the S H in the proline molecule, (CH2)3(CHCOOH)KH,lomers the basicity of the nitrogen. Thus, pKz for proline is 10.640 a t 2502 as compared with 11.305 for pyrrolidine. I t is well known that a hydroxyl group, by virtue of its electron-attracting nature, also lowers the basicity of nearbg nitrogen atoms. The value of pK2 for hydroxyproline is 9.662 at 25°.2 The elucidation of the effect of structure of acids and bases upon the change of ASo and ACPofor dissociation processes is complicated by the profound effect of electrostatic terms arising from differences in charge type. Whereas the dissociation of pyrrolidinium ion is an isoelectric process, BH+ = B -t H+, both of the dissociation steps of proline and hydroxyproline involve hybrid ions or zwitterions (designated by the superscript r ). The process usually designated as the “first dissociation” is the dissociation of the carboxyl (16) L. Sacconi, P. Paoletti, and M. Ciappolini, J . A m . Chem. Soc.. 82, 3831 (1960). (17) The negative R i m ascribed t o AS for the neutralization reaction appears to he i n error.
11:27
+
group, which can be represented BHH + = BHT H +. At high pH values, a proton is removed from the nitrogen (“second dissociation”), and the charge effects are asfollows: BHF = BII+. Seither of these is an isoelectric process. Although the behavior of the activity coefficient of a hybrid ion in solution has sometimes been found to approximate that of an uncharged molecule, it seems likely that the extent of solvation (upon which the inagnitudes of ASo and ACPoappear to depend) will often be quite didferent. If the two charge centers in the molecule B F interact with solvent dipoles independently,l8 it may be imagined that the second dissociation step matches more closely, from the electrostatic viewpoint, the dissociation of pyrrolidinium ion than does the finit. If, on the contrar-y, “overlapping” of the charges in the hybrid ion creates virtually an uncharged species, the first dissociation may be expected to approximate an isoelectric process. The changes of entropy and heat capacity for the two steps in the dissociation of proline and hydroxyproline have been recalculated from the data of Smith, Gorham, and The values at 125’ are as follows:
+
J.
First dissociation Proline Hydroxyproline Second dissociation Proline Hydroxyproline
ASO, deg.? mole-’
A Cpa j. des.-’ mole I
-32 1 -22.0
- 184 - 123
-60 6 -53 8
0 - 51
-1
From a comparison of the entropies, it would appear that the first, dissociation of proline parallels moire closely the dissociation of pyrrolidinium ion than does the second. The large negative change in heat CBpacity is, however, not a usual accompaniment of truly isoelectric dissociation processes1Qmost of which, like that for pyrrolidinium ion, have positive values of AC,O. As molecular models show, pyrrolidine can be formed from diethylamine with little distortion of the bonds, and a comparison of the thermodynamic properties of these two bases is therefore pertinent. Fyfe20found ~ J Y 10.98 for diethylamine at 2 5 O , in fair agreement with the value of 10.92 interpolated from the data of Evans and Hainann.21 For AHo, Fyfe found 43.1 kj. mo1e-l whereas Evans and Hamann give 53.4 kj. mole-l. The corresponding entropy values are -66 and -30 j . deg.-l mole-l, respectively. I n view of these differences, it is difficult to pursue the comparison with pyrrolidine. On the assumption that the thermodynamic properties of diethylamine and pyrrolidine should not differ greatly, however, the data of Evans and Hamann might perhaps be preferred. (18) See, for example, X. Bjerrum, 2. phusik. Chem., 104, 147 (1923); I. AI. Kolthoff and L. S. Gum, J . Am. Chem. Soc., 60, 2516 (1938). (19) See Table I11 of ref. 9. ( 2 0 ) W. S. Fyfe, J . Chem. Soc.. 1347 (1965). (21) A. G. Evans and S. D. Harnann, Trans. Faraday Soc., 47, 34 (1951).
