NOTES
would appear lines from cells adjacent to hydrogencontaining cells. Thus a composite broad line would occur, and the shape of the peaks produced by a recording densitometer would indicate how many of the possible spaces were filled with hydrogen. Such a search has been made in this laboratory,5 comparing X-ray diffraction films of vanadium and vanadium hydride. Within experimental error, the line widths of hydride and metal are identical, and thus it is unlikely that mechanical distension of the type described is occurring.
Acknowledgment. The authors are indebted to the U. S.Atomic Energy Commission for financial support for this work. (5) J. Macmillan, private communication.
Complexes of Azanaphthalenes with Iodine’
by Ivor Ilmet and Myron Krasij Department of Chemistry, University of Connecticut, Storrs, Connecticut (Received J u l y !2i9 1986)
This study was undertaken to test further the general validity of the relationship between the pKa of a heterocyclic aromatic amine and the log K of its iodine complex, as demonstrated by Chaudhuri and B a q 2 as well as to investigate the variation in the wavelength of the blue-shifted iodine absorption band with the pK, of the donor in azine-iodine complexes. While a linear dependence of the wavelength of the blue-shifted iodine band on the n-ionization potential of the donor has been demonstrated for complexes with such donors as halogenated hydrocarbons,a no such relationship seems to have been observed for azine-iodine complexe~.~Krishna and Chowdhuri5 have shown that the wavelength of iodine absorption does decrease with increasing pKa of the donor for a series of azine-iodine complexes, but owing to the small numbers of donors included in their study together with considerable scatter of the experimental points, no conclusion could be drawn as to the exact nature of the relationship. Complexes with quinoline and isoquinoline were included in the present study as there was some question as to the validity of previous results.2
Experimental Section Nerck’s Spectroquality chloroform was passed through a basic alumina column under nitrogen at-
3755
mosphere just before use. Quinoline (Fisher), isoquinoline (Eastman), and quinaldine (Aldrich) were dried over potassium hydroxide and distilled in vacuo. Iodine (Fisher), phthalazine, and l,5-naphthyridine (both Aldrich) were sublimed, and cinnoline (Aldrich) and quinoxaline (Eastman) were distilled in vacuo. Quinazoline was prepared by a known method.6 Solutions of complexes were prepared immediately (less than 5 min in all cases) before measurement from freshly prepared stock solutions of iodine and the donors by means of calibrated volumetric ware. All solutions were prepared and transferred under nitrogen atmosphere and were kept in dark bottles under nitrogen. The spectra were measured with a Cary 14 spectrophotometer with the cell compartment main0.1”. Nominally matched quartz tained at 25 cells were further calibrated by means of standard cupric sulfate solutions.’
*
Results and Discussion The wavelengths of the iodine absorption maxima and the molar absorptivities and formation constants as calculated by the Benesi-Hildebrand equationE are listed in Table I. The Benesi-Hildebrand relationship was used as there was, with one exception, negligible interference from adjacent absorption bands. T* Overlap was observed from the low-energy n transition of cinnoline in the cinnoline-iodine complex. The overlap could not be compensated for by an equal amount of cinnoline in the reference cell because the absorptivity of cinnoline in the complex appears somewhat enhanced over that of free cinnoline. Thus the errors in the molar absorptivity and the formation constant for the cinnoline complex are rather larger than for the other complexes. However, we for the do believe that the error in determining ,,,A, cinnoline complex is of no greater magnitude than for the other complexes. Representative spectra are shown in Figure 1.
-
(1) Abstracted in part from a thesis of If. Krasij, submitted in partial fulfillment of the requirements for the degree of Master of Science at the University of Connecticut, 1966. ( 2 ) J. N. Chaudhuri and S. Basu, Trans. Faraday Soc., 5 5 , 898 (1959).
(3) J. Walkley, D. N. Glen, and J. H. Hildebrand, J . Chem. P h y s , 33, 621 (1960). (4) See G. Briegleb, “Elektronen-Donator-Acceptor-Komplexe,”
Springer-Verlag, Berlin, 1961. (5) V. G. Krishna and 31. Chowdhuri, J . P h y s . Chem., 67, 1067 (1963). (6) If.T. Bogert and E. 31. McColm, J . Am. Ciiem. Soc., 49, 2650 (1927). (7) K. S.Gibson, National Bureau of Standards Circular 484, U. S. Government Printing Office, Washington, D. C., 1949. (8) H. A. Benesi and J. H. Hildebrand, J . Am. Chem. Soc., 71, 2703 (1949).
Volume 70,iV?tmber 11
Socember 1966
3756
NOTES
~~
Table I: Summary of Results of Azanaphthalene-Iodine Complexes in Chloroform
0.8
0.7
Donor
1. Isoquinoline
2. 3. 4. 5. 6. 7. 8.
0.6
& 0.5
H
f 8
2 0.4
Quinoline Cinnoline Phthalazine 1,SNaphthyridine Quinaldine Quinazoline Quinoxaline
393 398 401 409 413 417 420 430
81.2 64.8 36.6 25.0 20.7 12.5 10.8 5.90
2280 1520 3240 1960 1140 1920 1830 1500
5.40 4.94 2.42 3.47 2.91 5.41 1.51 0.56
Taken from A. Albert in “Physical Met,hods in Heterocyclic Chemistry,” A. R. Katritzky, Ed., Academic Press Inc., New York, N. Y., 1963.
0.3
0.2
6.0
0.1
5.0
4.0 350
400
450
Wavelength,
500
550
r
i1-
mp.
Figure 1. Representative spectra of azanaphthalene-iodine complexes in chloroform solution: (1) 0.0990 F isoquinoline, 0.000197 F iodine, 2-cm cell; (2) 0.0882 F quinoline, 0.000197 F iodine, 2-cm cell; (3) 0.0105 F cinnoline, 0.0350 F iodine 0.2-mm cell; (4)0.0374 F phthalazine, 0.000197 F iodine, 2-cm cell; (5) 0.0856 F 1,5-naphthyridine, 0.000197 F iodine, 2-cm cell; (6) 0.739 F quinaldine, 0.000197 F iodine, 2-cm cell; (7) 0.186 F quinazoline, 0.00500 F iodine, 0.5-mm cell; (8) 0.873 F quinoxaline, 0.000197 F iodine, 2-cm cell.
1.5
1.0
Log
With the exception of cinnoline and quinaldine complexes, a plot of pK, of the donor us. log K, shown in Figure 2, appears linear. I n the quinaldine complex the adjacent methyl group would be expected to interfere more with the bulkier iodine molecule than with a hydrogen ion. Therefore, a relatively weaker complex than expected from the pKa should result. I n case of the cinnoline complex we believe that the uncertainty in the determination of log K is not sufficient to account for the observed deviation, especially since a similar deviation is observed for of iodine plot, not shown, cinnoline in the pKa us. A,, despite a smaller error involved in the determination of. , , ,X The unusually low-energy n a* transition in cinnoline has been interpreted as arising from an interaction of the nonbonding orbitals of the vicinal nitrogens with the resultant formation of one nonbonding level of relatively high energy and one of low --)-
The Journal of Physical Chemistry
2.0
K.
Figure 2. A plot of the logarithm of the formation constant of the iodine complexes va. pK, of the donors. Numbering of the donors corresponds to that in Table I.
energy.9 As removal of this interaction would be a prerequisite for protonation but not necessarily for the formation of an iodine complex, a complex stronger than predicted from the pK, of the donor could conceivably result. This is in accord with the observation that in a plot of Amax us. log K , shown in Figure 3, cinnoline falls nearly on line. No similar effect is observed with phthalazine presumably because the nitrogen-nitrogen interaction in phthalaxine is much less than in cinnolineg so that any small effect could be overshadowed by experimental errors. The previously reported2 A,, and log K for the (9)
L.Goodman, J . Mol. Spectry.,
6 , 109 (1961).
KOTES
3757
430
420
E
-
i
8
410
400
390 1.5
1.0
2.0
Log K .
Figure 3. A plot of the logarithm of t h e formation constant of the iodine complexes us. the blue-shifted iodine absorption maxima. T h e numbering of donors follows t h a t of Table I.
quinoline-iodine complex appear to be in error. Since we were able to approximate the previous results by addition of small amounts of water or simply keeping the solutions open to atmosphere for a few days, we assume that the errors were due to impurities, mainly water. It appears that when applied with precaution and under rigorous control of experimental conditions such as exclusion of water vapor, A, of the iodine complex of a heterocyclic amine can be used to estimate its pK, in cases where other methods fail. For instance, the true pK, of quinazoline has been determined only recently'o because of the difficulties resulting from the presence of covalent hydration in water solution. A pK, of 1.56 could be estimated for quinaxoline from the present data which compares with the most recent value of 1.5.
Acknowledgment. This work was supported in part by a grant from the Research Foundation of the University of Connecticut. (10) Tlr. Armareggo, J. Chem. Soc., 661 (1962).
High-Temperature Enthalpy Studies of
thalpy difference, H , - HzeOOK, was measured for Bi2S3 and SbzSes by the drop-calorimeter technique. This enthalpy difference was measured in both the solid and liquid states so that the heats and entropies of fusion of these compounds could be determined. I n addition, the enthalpy difference for Bi2Sawas extended from 400' to the melting point at 763' so that the heat capacity of Bi& could be determined over this temperature range. The only previously reported determinations of the heat of fusion and high-temperature heat capacity for BizS3 have been reported by Kelley. Kelley3 determined the heat of fusion cyroscopically and obtained a value of 8.9 kcal/mole (17.3 cal/g). He4estimated the heat capacity from free energy data and obtained a value of 28.9 6.10 X loF3!!' cal/mole deg for the temperature range 298-1000OK. In addition, Romanovskii and Tarasov5 reported low-temperature heat capacity studies (65-300'K) on a relatively impure sample of Bi&. No measurements of these thermal properties for Sb2Se3have been reported.
+
Experimental Section Calorimeter and Accessories. Measurements were made using a drop calorimeter similar to that described by Goodkin, et aL2 The temperature rise of the calorimeter was measured automatically and printed out with a digital recorder. This arrangement of electronic equipment for measuring the temperature rise consisted of a Kiethley 149 millimicrovoltmeter, which measured the emf of a copper-constantan thermocouple in the calorimeter. The output from this voltmeter was converted to a frequency by a Dymec 2210 voltage-to-frequency converter, counted with a Hewlett-Packard 52313 counter and printed out every 20 sec with a Hewlett-Packard 566A digital recorder. Materials. The samples of Bi2S3 and Sb2Se3were prepared using semiconductor grades of the elemental constituents. The sample of BizSa was synthesized using the technique described by Glatz and Meikleham6 and the sample of Sb2Se3was synthesized in a similar manner. Samples of semiconductor grade copper and silver were also employed to determine the thermal ca-
Bismuth Trisulfide and Antimony Triselenide
by Alfred C. Glatzl and Karen E. Cordo Research Diviuion, Carrier Corporation, Swacuse, N e w York (Rereiced J u n e 16, 1966)
I n the course of investigating the physical properties of certain semiconductors in this laboratory, the en-
(1) NASA Electronics Research Center, Cambridge, Mass. 02139. (2) J. Goodkin, C. Solomons, and G. Jann, Rtv. Sci. Instr., 2 9 , 105 (1958). (3) K. K. Kelley, U. S. Bureau of Mines Bulletin No. 393, U. S. Government Printing Office, Washington, D. C.. 1936. (4) K. K. Kelley, U. S. Bureau of Mines Bulletin No. 406, U. S. Government Printing Office, Washington, D. C., 1937. (5) V. A. Romanovskii and V. V. Tarasov, Soviet Phys.-Solid State, 2 , 1170 (1960). (G) A. C. Glats and V. F. Meikleham, J . Electrochem. Soc., 110, 1231 (1963).
Volume 70, Number I 1
November 1966
