NOTES
335
Table I : Heat of Decomposition of S4N4
Table 11: Heat of Decomposition of SerN4"
m
(in uacuo),
OR,
-AEtat,
AEign.
g.
ohm
cal.
cal.
0.0038310 0.0038669 0.0035172 0.0036097 0.0038072
310.01 312.92 284.62 292.10 308.09
2.08 1.62 1.68 1.88 2.54
Run
6 7 8 9 10
0.51725 0.51791 0.46528 0.48342 0.50166
AEd
=
-111.16 f 0.47 kcal./mole
AEd"
=
-111.16 f 0.47kcal./mole
-A&
cal.
g.-1
595.73 601.44 608.58 600.77 609.51 Av. 603.21
Devistion, cal. g.-1
-7.48 -1.77 5.37 -2.44 6.30 f2.58
AHd' = -110.01 f0.47 kcal./mole AHf' = +110.01 f 0.47 kcal. /mole
m
(in vacuo);
Run
1 2 4 5 6 7 8 9
g.
0.05921 0.03281 0.G6737 0.08527 0.13822 0.14152 0.12406 0.11430
AR?
ohm
0.0002493 O.OOO1520 0.0002860 0.0004773 0.0006414 0.0007422 0.0006818 0.0006411
AEd = -148.86 f 5.97
cal.
20.32 12.39 23.32 38.91 52.29 60.50 55.58 52.26
-AEd, cal.
g. -l
Deviation, cal. g. -1
0.61 0.56 0.66 0.61 0.71 0.56 0.60 1.57
372.59 -27.71 360.67 -39.63 336.40 -63.90 449.21 48.91 373.23 -37.07 423.60 23.30 443.24 42.94 443.56 43.26 Av. 400.31 f 16.06 = -147.68 f 5.97
kcal./mole
AHt"
kcal./mole
The Heat of Formation of Se4Nl Samples of crystalline Se4N4were syntheslzed a t the University of Arkansas, shipped with caution to the University of Wisconsin, and decomposed in the bomb calorimeter by the same technique as used for S4N4. The elements were the only decomposition products and thus, by averaging the data from Table 11,one finds the heat of formation of solid Se4S4to be +147.7 f 6.0 kcal./mole. The samples were analyzed to be -99% pure, but the properties of Se4N4 (it explodes on the slightest touch when dry, much like NII) preclude excessive handling and purification. If anything, the most exothermic decompositions are probably the most characteristic since small amounts of decomposition may have occurred in the time between the synthesls and the calorinietric run for a particular sample. By using only runs 5 , 7 , 8 , and 9, which are most exothermfc and most consistent, one calculates AHrO = +163 f 3 kcal./mole and this latter value 1s recommended for Se&4(s). From this latter heat of formation, the standard heats of formation of Se(g)' and of S ( g ) , 4an estimated heat of sublimation for Se4S4of 20 f 10 kcal./mole, and assuming that Se4K4(s)has approxlmately the same
AEipn,
cal.
AHd'
kcal. /mole
AEd' = -148.86 f 5.97
S4N4,one calculates the average N-S bond strength in S4N4to be 73.5 * 1 kcal./mole. This number is consistent with the value for D(NS) = 115 =t25 kcal.,/mole on the argument that the bond in NS(g) is of order 2.5, while t h a t in S4N4is of order 1.65.6 If a direct proportionality were applicable, the average (S-S) bond in S4?';4(g) would be estimated as 75 kcal./ mole. With this model one concludes that the actual energy effects to be associated with S-S or N-N interactions in S4N4are small ( 5 2 0 kcal./mole). One further predicts the average N-S single-bond energy to be -45 kcal./moL:.
-AEt,t,
=
+147.68 f 5.97
kcal./mole
' Runs 1 and 2 were made on the original samples received; runs 4, 5, and 6 were made on a second group of samples; runs 7,8, and 9 were made on the last group of samples received.
structure as S4N4(s) with only N-Se bonds, one calculates the average N-Se bond energy in Se4N4(g)to be 59 f 10 kcal./mole. If the bond order is 1.65, by analogy with S4N4, then one predicts that D(N-Se) = 90 f 20 kcal./mole (bond order = 2.5); D(N-Se+) = 108 f 20 kcal./mole (bond order = 3.0) ; and that the average X-Se single bond energy rn 40 kcal./mole. Acknowledgment. The authors are pleased to acknowledge the support of this work by the United States Atomic Energy Commission, by the American Chemical Society through a grant from the Petroleum Research Fund, and by the Selenium-Tellurium Development Association. (6) D. Chapman and T. Waddington, Trans. Faraday Soc., 5 8 , 1291, 1679 (1962).
(7) G. N. Lewis, M. Randall, K. S. Pitzer, and L. Brewer, "Thermodynamics," 2nd Ed., Mc Graw-Hill Book Co., New York, N. Y., 1961.
Association of Secondary Amines w i t h Tetrahydrofuranla
by H. Hartig and W. W. Brandtlb Department of ChJ.mistry,Illinois Znatitute of Technology, Chicago, Illinois 60616 (Received April 21, f964)
The steric effects modifying the H-bonding tendency of certain polar groups are of great interest because they may well determine certain chemical and physical Volume 69, Number 1
January 1966
KOTES
336
rate constants. Smith and Creitz2 obtained infrared data showing that 3-pentanols with increasing numbers of alkyl groups close t'o the hydroxyl group are incapable of forming H-bonded polymers or even dimers. Similarly, Bellarny and Williams3 studied several phenols and found the differences between the OH stretching frequencies of t'he H-bonded and nonbonded state to be relat,ively small whenever the OH group was strongly sterically hindered. Earlier studies in this laboratory4 showed that' the carbonyl group's ability to form H bonds can be markedly reduced by the presence of bulky groups on neighboring carbon atoms. The present study is concerned with the >N-H group, which is of great interest in the study of polyamides, proteins, and polypeptides. Bellaniy and Williams3 noticed t>hatthe constants of the H-bonding equilibrium were much more sensitive to the presence of sterically hindering groups than the difference bettween the frequencies of the absorption bands of the nonbonded and the bonded species, respectively. These authors concluded that in t'he series of compounds studied, the steric effects are primarily affecting the ease of H-bond formation, and only secondarily the bond strength. Bot,h frequency shifts and equilibrium constan& of H-bond format,ion were determined in the present st,udy, so as to see if the findings of Bellaniy and Williams hold for secondary amines as well. The choice of secondary amines for comparative studies is somewhat' limited by difficulties in their synthesis or purification and by their limited stability and solubility in interest'ing solvents. Tetrahydrofuran is a convenient, proton accept,or because it.s infrared and ultraviolet spectra do not interfere with those of the secondary ainines, it is unreact'ive under t,he experimental conditions, and it is a good solvent for the amines chosen. Also, there are some literature data on diphenylamine in tetrahydrofuran5 which may lead t.0 important comparisons. The problem of obtaining accurate constant's for association equilibria has been discussed in some detail by other workers.Ha'b To gain confidence in t'he result's of this work two spect'rophotometric methods (infrared and ultraviolet) were used. X.1n.r. chemical shift ineasurenients of the amino hydrogen signal were found to he riot feasible with the available equipment.
used. The di-o-tolylamine was prepared by the third procedure of Weston and Adkins7 with copper dust prepared from zinc and copper sulfate as catalyst. The compound was recrystallized several times, the last time from 2,2,4-trimethylpentane. The ni.p. was 51.5-52.5' before the last step (lit.8 52-53). Dip-tolylamine (E(and K Laboratories) was also recrystallized from the same solvent until the ultraviolet spectrum remained constant and showed no shoulder a t 2850 8. Diphenylamine (Purissimurn grade of Fluka A. G., Switzerland) was recrystallized once from 2,2,4-trimethylpentane. Di-2,2-naphthylamine (-41drich Chemical Co.) also was recrystallized from the same solvent until its ultraviolet spectruiii became constant. Di-1 ,1-naphthylamine was prepared by the method of Hodgeon and Afersdeng and was recrystallized until its melting point remained constant at 115116.5'. The solutions needed were prepared by drybox techniques and the exposure to the atmosphere during the spectroscopic measurement was kept to a minimum because of the known hygroscopicity of tetrahydrofuran. To avoid systematic errors, all solutions and reference solvents in a given series of runs were treated in a similar fashion. ( B ) Infrared Measurements. A Perkin-Elmer Uodel 21 spectrophotometer with a NaCl prisni was used. The slit width was 50 nip and it was found that the transmittance measured at half this slit width was lower by 3%. Froin the research of Philpotts and co-workers,'O it appears that in the present experiments the measured peak heights are all affected to the same fractional extent because the infrared absorption bands investigated in the present work are broader and the slit widths lower than those used by these authors. I t can be shown that the equilibriuni constants do not carry systematic errors from this source.
Experimental Methods and Procedures (4) Materials. The hydrocarbon solvents used were Spectroquality reagent grade from Matheson Coleman and Bell, Inc., or from the Phillips Petroleum Co. Tetrahydrofuran (THF) was Fisher Certified grade. The solvents were redistilled over sodium and tmhecenter fraction distilling within a range of 0.3' was
(1963).
The Journal of Physical Chemistry
(1) (a) This work was supported by a grant from the U. S.Public Health Service (GM-10288); (b) direct requests for reprints to this author. (2) F. A. Smith and E. C. Creitz, J . Res. S a i l . B u r . Std., 46, 145 (1951). (3) L. J. Bellamy and R. L. Williams, Proc. Roy. SOC. (London), A254, 119 (1960). (4) W.W. Brandt, J . Am. Chem. Soc., 85, 2628 (1963). (5) A. B. Sannigrahi and A. K. Chandra, J . Phys. Chem., 67, 1106 (6) (a) N. J. Rose and R. S. Drago, J . Am. Chem. SOC.,81, 6138 (1959); (b) P. R. Hammond, J . Chem. SOC.,479 (1964). (7) 1'. E. Weston and H . Adkins, J . Am. Chem. SOC.,50, 859 (1928).
(8) I. M. Heilbron, et al., "Dictionary of Organic Compounds," Oxford University Press, New York, N. T.,1934. (9) H. Hodgeon and E. Mersden, J . Chem. Soc., 1181 (1938). (10) A. R. Philpotts, W. Thain. and P. G. Smith, Anal. Chem., 23, 268 (1951).
NOTES
337
Table I : Infrared Measurements on Solutions of Secondary Amines in Mixtures of Tetrahydrofuran ( T H F ) and Hydrocarbons* Amine
r
Concentration, niole/l. Hydrocarbon solvent Peak positions, cm. -l Free >N-H Bonded >S-€1 in hydrocarbon in tetrahydrofuran Molar extinction coefficients, l./mole cm. Free >N-H in hydrocarbon Bonded >S-H in tetrahydrofuran K , l./mole
0.02 0.2 2,2,4Trimethylpentane
-
Diphenylamine
Di-o-tolylamine
0.2 Cyclohexane
0.2 2,2,4-Trimethylpentane 3442
3459
3465
(3380) 3364
(3380) 3362
(3378) 3363
(3386) 3374
29 113 1.53 & 0.15
29 119 1.45 f 0 . 1
118 1.53 f 0 . 1
15 56 0.44 f 0.02
T H F concentrations range from 0.06 to 12.2 moles/l.; temperature, 22". Values in parentheses are extrapolated to infinite T H F concentration. K equilibrium constant evaluated by the method of Rose and Dragee" a t the peak maxima of the bonded >N-H bands.
( C ) Ultraviolet Measurements. A Cary Model 14 spectrophotometer of the Applied Physics Corp. was used, using a slit width of about 0.4 mm. It was shown experimentally that the peak heights are not dependent on slit widths under the experimental conditions. ( D ) Data Treatment. The ultraviolet molar extinction coefficients of the amines (or diphenyl ether) in a series of solutions containing various amounts of tetrahydrofuran (THF) were used to obtain K values from the equation" (see Figure 1)
Here [BIo is the analytical concentration of the H acceptor B, while E, eo, and el are the apparent molar extinction coefficient of the proton donor, the actual extinction coefficient of the unbonded proton donor, AH, and of the bonded proton donor, A H - - - B , respectively. The equation is valid if [AH - - - B ]
