NOTES
for BrgO. We found that Xilman’sll application of the Estrup-Wolfgang kinetic model**-c to the liquid bromine-cyclohexane system, assuming hot bromine atom reactions with cyclohexane molecules, did not correlate with our zero-time bromine data. I n fact, the linearity of the scavenger curve supports a radical-type modelsc such as the random fragmentation mode1,12 the Auger or a combination of both. radiation The Auger radiation hypothesis was given increased support by the work of Thompson and Miller13which shows that in condensed media, certain recoiling atoms rarely develop a charge prior to loss of initial kinetic energy. If Auger radiation alone were responsible for the final chemical state of the BrBDm and Brsz one might expect an isotope effect due to a different per cent of internal conversion. l 4 However, if Auger radiation led to the same per cent organic combination as random fragmentation, no isotope effect would be expected. Since none was observed (in the absence of postactivation thermal reactions), the Braomand BrB2 hot-atom yields in Brz-scavenged cyclohexane may be explained by the random fragmentation model or a combination of random fragmentation and Auger radiation models. The nature of the reactive species, its control, and the possible role of BrBZm are being extensively studied. If a long-lived, thermally reactive species is formed as a consequence of isomeric transition reactions as well as by (n,r) reactions, the observed pseudo-isotope effect may be due to an increase in the concentration of the species formed by the Br82misomeric transition reaction. Preliminary results indicate that such species may be formed in n-hexane, 2,2-dimethylbutane1 2,3dimethylbutane, 2-methylpentane, and 3-methylpentane. Ellgren15 has found postactivation thermal reactions of Br2 in CC1, with an increase in the organic retentions of BrBomand Brs2with length of irradiation. If such long-lived reactive species are formed by radiobromine atoms in all organic solvents, it may call for a re-evaluation of most liquid bromine hotatom chemistry data. Acknowledgment. We thank Alan J. Blotcky and the personnel of the Omaha V.A. Hospital who assisted us in the many reactor irradiations. This research was supported in part by funds made available by the University of Nebraska Research Council.
(12) J. E. Willard, Ann. Rev. Nucl. Sci., 3, 193 (1953). (13) J. L. Thompson and W. W. Miller, J . Chem. Phye., 38, 2477 (1963). (14) S. Wexler and T. H. Davies, ibid., 20, 1688 (1952). (15) W. X. Ellgren, private communication, this laboratory, 1964.
2799
Hydrogen Bonding of Amide Groups in Dioxane Solution by H. Susi Eastern Regional Research Laboratory, Eastern Utilization Research and Development Division, U.S. Department of Agriculture, Agricultural Research Service, Philadelphia, Pennsylvania 19128 (Received February 1 , 1966)
Lactams frequently have been used as model compounds to investigate amide to amide hydrogen bonding because the cis arrangement of C=O and N-H groups leads to dimerization exclusively and prevents complications arising from multiple equilibria between hydrogen-bonded chain polymers. l-a Various authors have employed infrared methods to investigate the dimerization of lactams in carbon tetrachloride, a solvent which is virtually inert from the standpoint of hydrogen b 0 n d i n g . 1 ~ ~Relatively little work has been reported in solvents which are themselves capable of hydrogen-bond formation. Such studies lead to a net energy difference between initial and final states characterized by amide groups bonded to the solvent and to each other, respectively. The results are of interest in assaying the role of hydrogen bonds in the secondary structure of proteins under conditions where amide-amide bonds can compete with amide-solvent bonds. The energy of interaction of even weak proton donors, such as chloroform, with an amide has been estimated a t 2 kcal./mole.2 This communication reports a study of the dimerization of 8-valerolactam in a proton-accepting solvent , dioxane, by methods of infrared spectroscopy. One of the principal difficulties encountered in work with complex systems is proper evaluation of the temperature dependence of the absorption coefficients of various species.5 For the investigated system, a procedure involving some approximations is proposed which takes into account both solvent-solvent and solvent-solute interactions and which leads to internally consistent results over a wide range of temperature and concentration.
(1) M.Tsuboi, Bull. Chem. SOC.Japan, 24, 75 (1951). (2) W. Klemperer, M. W. Cronyn, A. H. Maki, and G. C. Pimentel, J . Am. Chem. SOC.,76, 5846 (1954). (3) R. C. Lord and T. J. Porro, 2. Elektrochem., 64, 672 (1960). (4) H. E. Affsprung, S. D. Christian, and J. D. Worley, Spectrochim. Acta, 20, 1415 (1964). (5) G. C. Pimentel and A. L. McClellan, “The Hydrogen Bond,” W. H. Freeman and Co., San Francisco, Calif., 1960, p. 77.
Volume 69, Number 8 August 1966
NOTES
2800
Experimental 8-Valerolactam (Aldrich Chemical CO.~) was purified by distillation to remove traces of water. Dioxane (Fisher Certified Reagent) was refluxed with dilute hydrochloric acid for 12 hr., then refluxed over sodium for 24 hr. and distilled.’ Absorption spectra from 6000 to 7000 cm:-’ were obtained with a Cary Model 14 instrument. Sample concentrations ranged from 0.1 to 0.8 mole,’l. The absorption of each sample was investigated a t 25, 45, and 65’. A water-jacketed controlled-temperature absorption cell of conventional design‘ and 25-mm. path length was employed. Only one temperature-dependent absorption band of appreciable intensity, centering close to 6690 cm.-l (-1.495 p ) , was observed in the investigated range. This band was assigned to the first overtone of the NH stretching mode of amide groups not involved in amideamide H-bonding, on the basis of its temperature dependence and by comparison with spectra obtained in various other solvents (CCl,, CC1,H) and the spectrum of the pure lactam. I t should be pointed out that neither in dioxane solution nor in any other investigated solvent (or in the spectrum of the pure lactam) were bands of appreciable intensity found which could be assigned to overtones of N H stretching fundamentals of groups involved in amide-amide bonding. Although the corresponding fundamentals are very intense, 1-3 the overtones must be extremely weak. Figure 1A gives the observed peak absorbance of the investigated absorption band as a function of temperature and concentration, as observed in a 25-mm. cell. Evaluation of Experimental Results
If no higher polymers are formed and if the solvent is inert, then the following relations hold
+ ca)
K’, =
(3)
where K‘, is the dimerization constant in mole fraction units; C is the total solute concentration in moles per liter; Cais the solvent concentration in moles per liter; M is monomer concentration in moles per liter; A is the absorbance of monomers; e is the absorptivity of monomeric species; and d is the path length. Rearrangement of (3) leads t o Ca/A
=
+
(Ed/A2)(C2/2 cat)
The unknowns K‘, and
E
The Journal of Physical Chemistry
-
(4K’z
+ 1)/2Ed
(4)
can be obtained by setting
and by plotting y 2)s. x. Ed is given by the slope and K‘, is calculated from the intercept. The necessity to evaluate E by extrapolation is eliminated.
200
Y 100 I50 50
/A:
0
B
/*
e* *e.,
-----------I
I
I
Figure 1. A: Observed absorbance of 8-valerolactam in dioxane solution a t 1.495 p ; path length: 25 mm. B: Grctphical evaluation of absorptivity and dimerization constant; y = C,/A; 5 = (1/A)2(C2/2 C,C).
+
Figure 1B presents the results of the described graphical procedure. I t is immediately obvious that the slope of the obtained lines varies with temperature, ie., that the observed absorptivity [eobad = (dA/dC)T,c,~] is strongly temperature dependent. An inspection of Figure 1A leads to the same conclusion. If the monomer is heavily bonded to the solvent, then dimerization occurs by breaking up this complex and forming the dimer and two solvent molecules. In this case, eq. 1 holds for the dimerization of initially solvent-bonded monomers. Previous infrared studies have shown that amides are very strongly bonded to ethers18ie., the monomer-solvent association constant (6) Mention of specific firms and products does not imply endorsement by the U. S.Department of Agriculture over others of a similar nature not named. (7) K. Hess and H. Frahm, Ber., 71, 2627 (1938). (8) S : Mizushima, M. Tsuboi, T. Shimanouchi, and Y . Tsuda, Spectrochzm. Acta, 7,100 (1955).
NOTES
2801
has a high value. If the association constant for solutesolvent interaction is much higher than the solventsolvent association constant, eq. 3 is also applicable to the dimerization of initially solvent-bonded monomers, provided the temperature dependence of the apparent absorptivity is taken into consideration. Figure 1B leads to the following numerical values: T = 298OK., Eobsd = 0.184 1. mole-l cm.-I, K', = 4.1 (mole fraction)-l; T = 318O, Eobad = 0.246, K', = 4.2; T = 338O, Eobsd = 0.295, K', = 4.0. The variation of K', is within experimental accuracy. Standard methods lead to: AH = 0 0.5 kcal./mole, AS' = 3 =t0.5 e.u. for the dimerization of initially solvent-bonded monomers. The entropy value refers t o a standard state defined in mole fractions.
+
Diwcussion The results suggest that the energy required to break amide-ether hydrogen bonds is approximately equal to the energy of formation of amideamide bonds. The relatively high value of the association constant is thus a consequence of a positive standard entropy change upon dimerization. While the system is far too complex for a detailed discussion of entropy relations, a brief qualitative consideration might provide some insight. The "monomeric" system involves two particles schematically represented by: 2 X (amidedioxane). Upon dimerization, one amide-amide dimer (with two H bonds) and two free solvent molecules are formed, i.e., the total number of particles is increased. At sufficiently high concentration, the positive value of ASo results in a considerable number of NH---O=C bonds being formed without a net change in enthalpy. The value of K', reported here is about 1/725 times the value given by Tsuboil for dimerization of 6-valerolactam in carbon tetrachloride solution (Kc(cc4, 26') = 280 moles/l.; Kz(Cc4,26") = 2900 (mole fraction) -l), reflecting the strong competition between dimerization and solvent interaction in the studied system. The experimental data, as presented in Figure lA, re-emphasize the necessity31sto determine the temperature dependence of apparent absorptivity values if thermodynamic quantities are deduced from absorbance measurements in complex systems. Our results also lend support to the conclusiongthat the anharmonicity of X H stretching fundamentals is reduced by hydrogen bonding.lO (9) See ref. 5, p. 114. (10) The overtone frequency in dioxane solution is very close to the overtone frequency in CCla solution. The fundamental frequency is substantially lower in dioxane solution.
Ion-Pair Formation in Aqueous Solutions
of Butylammonium Isobutyratel
by Irving M. Klotz and Henry A. DePhillips, Jr.2 Departmnt o j Chemistry, Northweatern University, Evanston, Illinois (Received February 8, 1966)
Among the noncovalent bonds which may participate in the fixation of conformation in protein molecules one must include the possibility of electrostatic interactions between pendant -NH3+ and -COO- side chains. From the physicochemical behavior of simple ions in aqueous solution there are indicationsa that ion-pair formation, such as [-NH3+. .-OOC-], does not occur to any detectable extent except at very high concentrations. Nevertheless, no direct measurements have been available for a system of an ammonium and carboxylate group suitable as amodel for potential interactions among protein side chains. We have now found that optical measurements in the near-infrared, similar to those used previously to study amide hydrogen bonds,4 do give an insight into ion-pair formation between n-butylammonium and isobutyrate ions. In essence we have used the overtone infrared region to measure the equilibrium constant for the reaction
-
CdHsNH3+
+ C3H7COO[ C ~ H ~ N HC3H,COO] S, +*-
(1)
From the equilibrium constant, established a t three temperatures, the free energy, enthalpy, and entropy of ion-pair interactions may be calculated.
Experimental Preparation of Solutions. Spectra were examined in the region of 1.5 p rather than in the fundamental range so as to minimize complications due to the absorption by water, the solvent. As in previous work a reference cell was prepared containing the same amount of water as the sample cell but with a solute, (1) This investigation was supported in part by a grant (GM-09280) from the National Institute of General Medical Sciences, Public Health Service. It was also assisted by support made available by a Public Health Service Training Grant (No. 5T1-GM-626) from the National Institute of General Medical Sciences. (2) Predoctoral Fellow of the U. S. Public Health Service, 19611963. (3) I. M.Klotz, Brookhaven S y n p . Biol., 13, 26 (1960). (4) I. M. Elotz and J. S. Franzen, J. Am. Chem. Soc., 84, 3461 (1962).
Volume 69, Number 8 August 1966
