HARID. BISTAND WILLISB. PERSON
3288
Thermodynamic Properties and Ultraviolet Spectra of Cyanogen Iodide Complexes with Some n Donors
by Hari D. Bistl and Willis B. Person2 Department of Chemistry, University of Florida, Gainesuille, Florida
(Received January 87,1967)
The ultraviolet absorption spectrum of I C N in n-heptane and of its complexes with diethyl ether, ethanol, and tetrahydrofuran has been investigated in n-heptane solution. From the “blue-shifted” ICN band a t 240 mp, the thermodynamic properties K ,AH’, and ASo and the characteristics of this absorption band (emax, A v ~ , f, ~ , and D)were determined for the complexes. Comparison of these properties with the corresponding values for IZ complexes establishes the similarity in both thermodynamic and spectroscopic behavior. The failure to observe the charge-transfer absorption is discussed. It is concluded that this failure is due t o a shift of h v C T to about 1-2 ev of higher energy (as compared to I2 complexes) because of a smaller vertical electron affinity (estimated to be only 0.9 ev for ICN) and a more negative value of Go for the ICN complexes.
Introduction In spite of its potential interest as an electron acceptor to be compared with 12,ICl, and other halogens, very little information exists about electron-donor-acceptor complexes with ICN. Fairbrother3 measured the apparent dipole moments of ICN in different solvents, noting an appreciable increase for p in solutions with electron donors as solvents, similar to that which he had found earlier for 12.4 In that paper3 he commented on and reported experiments confirming the frequency shifts in the ultraviolet spectrum of IC?; observed earlier by Gillame5 Further studies of the changes in the ultraviolet spectra in complexing solvents are reported by Haszeldine,6 and changes in the infrared spectra were reported by Haszeldine,8 by Glusker and Thompson,’ and by Person, Humphrey, and Popov.* Formation constants for a few complexes have been reported from measurements made in the infrared r e g i ~ n . ~However, only a relatively few complexes have been studied even that quantitatively. One question arising from these studies concerns the frequency of the charge-transfer absorption band which presumably exists for these complexes. Haszeldine6 reports some changes in the ultraviolet spectrum near the limit of transmission of his spectrometer-Le., around 200 mp. Badger and Woo’O and Yakovlevall have reported the spectrum of ICN vapor in this region. The Journal of Physical Chemistry
Badger and Woo’O assign the broad absorption found here (with peak a t 247 mp) to a transition to an electronic upper state which dissociates to give the products I(2Pa/,) CN*(2TI). This transition then corresponds roughly to the transition for the visible absorption band of I2 (mostly to the upper state, which dissociates to give I(2P8,z) I*(2Pl,,). Hence, one might expect the absorption in this region for the u acceptor, ICN, to exhibit a change on complexing similar to that found for the visible I2 band-ie., a “blue shift” and intensification.12 The possibility that the charge-transfer ab-
+
+
~
(1) Fulbright Exchange Fellow, 1963-1965. (2) Visiting Associate Professor of Chemistry and National Science Foundation Senior Postdoctoral Fellow, Laboratory of Molecular Structure and Spectra, Department of Physics, University of Chicago, Chicago, Ill., 1965-1966. The experimental work was done at the University of Iowa. (3) F. Fairbrother, J. Chem. SOC.,180 (1950). (4) F . Fairbrother, ibid., 1051 (1948). (5) A. E. Gillam, Trans. Faraday SOC.,29, 1132 (1933). (6) R . N. Hasaeldine, J . Chem. SOC.,4145 (1954). (7) D. L. Glusker and H. W. Thompson, ibid., 471 (1955). (8) W. B. Person, R. E. Humphrey, and A. I. Popov, J. A m . Chem. SOC.,81, 273 (1959). (9) A. I. Popov, R. E. Humphrey, and UT.B. Person, ibid.,82, 1850 (1960). (10) R. M. Badger and S. Woo, ibid., 53, 2572 (1931). (11) A. V. Yakovleva, Izv. A M . Nauk SSSR, Ser. Fiz., 14, 517 (1950); Chem. Abstr., 45, 3239b (1951).
CYANOGEN IODIDE COMPLEXES WITH SOME n DONORS
sorption band of complexed ICN is superimposed on the absorption in this region by the free ICN should not be overlooked, however. Because the changes in the ultraviolet region of the spectrum of complexed ICN do occur in the region below 250 mp, the choice of donors which can be studied is limited. The requirement of transparency rules out aromatic a donors, such as benzene, as well as many interesting n donors-e.g., sulfides. The requirement of stability rules out the stronger n donors-e.g., trimethylamine-which react further with ICN. Hence, only the ethers and alcohols, among the usual kinds of donors, form complexes which can be studied in the ultraviolet region. We report a study of the complexes of ICN with diethyl ether and with ethyl alcohol, in an attempt to obtain the thermodynamic properties K , A H , and AS for complex formation, and to obtain further information about the spectral changes of ICN near 250 mfi. I n an attempt to provide some information about the interaction of ICN with a stronger n donor, we report also some much less complete results of a study of the complex with tetrahydrofuran.
Experimental Section Chemicals. Iodine cyanide was prepared by the method of Bak and HilleberV and was purified by recrystallizations from chloroform. Solvents and donors were purified according to standard procedure^.'^-'^ n-Heptane (Phillips Petroleum Co., pure grade) was shaken with a mixture of concentrated HNO, and HzS04 and again repeatedly with concentrated H2S04,washed, and distilled under nitrogen over sodium. Ethyl alcohol (95%) was refluxed over freshly burnt quicklime and then distilled. “Superdry” alcohol was prepared from this by distilling it again from magnesium activated with Iz. Ether was stored over sodium and fractionally distilled. The tetrahydrofuran was purified by removing peroxides and dried well, immediately before use. The solutions were freshly prepared on the day of use. The donor and ICN were weighed into a 25-ml volumetric flask, with a semimicro Mettler balance, and the volume was adjusted with n-heptane. For reference in the double-beam spectrometer, solutions containing solvent and donor a t the same concentration as in the ICN sample solution were also prepared. Spectra. The spectra were recorded in the region 360-170 mp (or over portions of this region) with a Beckman DK-2A far-ultraviolet spectrophotometer. A matched pair of Beckman far-ultraviolet silica cells (1 cm) was used. These cells were mounted in the Beckman temperature-regulated cell holder, whose temperature was controlled by circulating a glycol-
3289
water mixture from a Haake thermostat and, for low temperatures, from a Sargent water bath cooler. The temperature in the cells could be controlled to *0.2” in the range from - 12 to 60”. Oxygen and other undesirable vapors were removed by purging the instrument with Linde high-purity dry nitrogen. A slow purge was maintained continuously, and the purge rate was increased during the study, especially for the low-temperature studies in order to prevent condensation of H20. The cells containing the ICN solution and the reference solution were placed in the spectrometer and the spectra were recorded a t 28”. Then the temperature was adjusted in steps and the spectra were recorded a t each temperature, after allowing 30 min or more for the cell compartment to reach equilibrium. Spectra were recorded in the order: 28, 19, 5, -10” (45 and 60”, in the case of the alcohol complex). Actually, the temperature of the cells was adjusted only approximately to these values. The spectra were read, and the absorbance was plotted as a function of temperature for each solution. The absorbances a t each of the temperatures above were read from the graph by interpolation. I n order to check that possible irreversible reactions did not occur during the time needed to go through this cycle, the spectrum of the solution a t 28” was usually recorded at the end of the cycle as well. Dissolved oxygen in n-heptane shows considerable absorption in this region of the ultraviolet spectrumespecially below 2100 A. Stringent precautions were taken to avoid it including: (i) distilling all chemicals under nitrogen atmosphere just before use; (ii) purging both the donor and the solvent with extra dry highpurity nitrogen just before preparing the solutions under N2atmosphere in a drybox; (iii) filling the Beckman stoppered quartz cells in the drybox; and (iv) keeping the spectrometer, and especially the cell compartment, purged continuously with dry nitrogen. We did not detect any complications due to the contact charge-transfer absorption of Oz. A further small error (about 301,) occurred in the molar concentrations of the solutions owing to the changes in volume as a result of the temperature changes. This effect was ignored, since it was small (12) For example, see R. S . Mulliken, J. Chim. P h y ~ . 20 , (1964); R.9. Mulliken, Rec. Trav. Chim., 75, 845 (1956). (13) B. Bak and A. Hillebert, Org. Syn., 32, 29 (1952). (14) K.B. Wiberg, “Laboratory Technique in Organic Chemistry,” McGraw-Hill Book Co., Inc., New York, N . Y.,1960. (15) A. Weissberger, Ed., “Technique of Organic Chemistry,” Vol. VII, Interscience Publishers, Inc., New York, N . Y.,1955. (16) A. I. Vogel, “Elementary Practical Organic Chemistry. Part I. Small Scale Preparations,” Longmans, Green and Co., Ltd., London, 1957.
Volume 71, Number 10 September 1967
HARID. BIST AND WILLISB. PERSON
3290
Axf = ~xfl[IC,Nl (2) The value of cAf in the solutions containing donor is assumed to be the same as that found for ICN alone in n-heptane. Hence, we may compute the absorbance of the complex corrected for that due to free ICY
Acx
205
%5(mr,
-,245
265
285
305
325
Figure 1. The ultraviolet absorption spectra of Et20-ICN solutions; solvent, n-heptane; 23’; I.O@cm cell. The ICN and Et20 concentrations in millimoles per iiter are: (1) 3.368 and 0.000; (2) 3.368 and 39.76; (3) 3.466 and 107.99; (4)3.499 and 338.3; ( 5 ) 3.629 and 1297.3; (6) 3.695 and 9626.3.
and since quantitative information about the density changes in our solvent system was not readily available.
Results Figure 1 presents representative spectra of solutions in n-heptane of approximately constant ICX concentration with increasing concentrations of ether. It illustrates the decrease in the intensity of the free ICN band a t 241 inp (in n-heptane) as the complex is formed, with the resulting increase in the band due to the complexed I C N at 225 mp. The isosbestic point indicates that only two species exist in these solutions. We note that the absorption by this shifted I C s band at 225 nip in the complex is overlapped by the stronger absorption near 200 mp. It is still not absolutely clear whether this strong band results from a shift of the absorption band already seen near 200 mp, for ICN in n-heptane, or whether a charge-transfer absorption band for the complex is superimposed on the ICN absorption in this region. However, we believe the arguments in favor of the former interpretation of our spectra are reasonably convincing. From a set of absorption curves such as those shown in Figure 1 for the solution at a given temperature, the formation constant and molar absorptivities can be determined in the usual way.’’ Since the absorption of the complex overlaps that of the free ICK, a correction must be made for the latter. The concentration of free ICK ia given by [XCS]
=
(1CN)o -
K (1CN)oDo 1 KDo
+
(1)
Here the aero subscript denotes the total concentration of t,he species, determined from the weight. The absorbance a t X due to free ICN is The Journal of Physical Chemistry
Ax - Axr
(3) Here A x is the total observed absorbance at X in the solution, and Acx is that due to complex alone. We may now use AQ, measured in a series of solutions with differing- donor concentrations, in the BenesiHildebrandls or Scottlg equations to obtain the value of tCX and K . (Or, better, we may use the Liptay20 procedure to treat these data*) However, in Order to make the corrections of eq 113, we must estimate K . We have done this using Nagakura’s equationz1 K =
=
+
Do(Ao - A ’ ) Do’(A - Ao) D,Do’(A’ - A )
(4
Here Ao, A , and A’ are the absorbances of solutions whose total donor concentrations are 0, Do, and Do’ moles/l., respectively. After making this estimate for K , the corrected absorbance is computed by eq 1-3 and then treated by the Liptay procedure20a22 to obtain K and tcXvalues, cycling until K does not change. The Liptay analysis of the data was programmed22 in FORTRAN IV for the IBM 7040/7044 computer. Absorbance data from five solutions with different donor concentrations were read at eleven wavelengths from 210 to 235 mp (at 2.5-mp intervals) for the ether complex, and at four wavelengths from 210 to 225 mp (at 5-mp intervals) for the alcohol complex. The resulting K values, together with tmsx values (at X 225 mp for the ether complex and a t X 217 mp for the alcohol complex) are tabulated in Table I. S o trends were noticed in the {matrix or in the D, matrix in the Liptay analysis,z0 suggesting that there is little noticeable effect owing to failure of either the assumption that exf is constant or that only 1: 1 complexes are present. This analysis was carried out at each of four temperatures from data obtained above and illustrated for a typical solution in Figure 2. This figure again illustrates the increase in absorption of the complex near (17) G. Briegleb, “Elektronen-Donator-Acceptor-Komplexe,” Springer-Verlag, Berlin, 1961, Chapter XII. (18) H. A. Benesi and J. H. Hildebrand, J . Am. Chem. SOC.,71, 2703 (1949). (19) R. L.Scott, Rec. Trav. Chim., 75, 787 (1956). (20) See ref 17; W. Liptay, Z . Elekfrochem.,65, 375 (1961). (21) S. Nagakura, J. Am. Chem. Soc., 80, 520 (1958), and other references cited there. (22) L. Julien, Ph.D. Thesis, University of Iowa, 1966.
CYANOGEN IODIDE COMPLEXES WITH SOMEn DONORS
3291
Table I : Equilibrium Constants, K , for Complex Formation, and emax Values for ICN Complexes with Tetrahydrofuran and with Diethyl Ether and Ethanol a t Different Temperatures
___Temp, OC
THF----
-9.5 5.0 18.0 27.0
13 1 4
---------
Diethyl ether-----K, I./mole ~~(225)
r - -
K, 1. /mole
fC(220)
17.1 f 1 . 9 8 . 3 f 1.0 6.2 10 . 7 3 . 8 i0 . 3
236 121
1 t
I OS8
e
255 f 7 260 3~ 14 241 f 15 257 f 9
Ethanol---------
Temp,
K,
O C
I. / m o 1e
t c (2 17)
-12.0 4.0 19.0 28.0
10.90 10 . 7 6.8 f 0 . 7 4.3 f 0 . 2 3.5 1 0 . 4
266 13 264 f 8 270 16 255 116
6.0
4.0
2.0
2 05
225
245
h
(md
-
265
285
305
Figure 2. Temperature dependence of the ultraviolet absorption spectra of an EtOH*ICN solution; solvent, n-heptane; 1.00-cm. cell. Concentrations for ICN and EtOH are 3.335 and 201.3 mM, respectively. The temperatures for the curves are: (1) 30.6, (2) - 12.2, (3) 5.8, (4) 19.2, (5) 46.8, (6) 59.8".
220 mp and decrease in absorption by the free ICN at 240 mp as more complex forms at the lower temperatures. Again, we note the presence of an isosbestic point.23 Wote that emax in Table I is independent of temperature within experimental error, as expected over this small range of temperature. We note also that the experimental uncertainty in K is rather high (f10%). However, we believe this uncertainty is acceptable for such weak complexes. We note that the results for C2H80H-ICN confirm the somewhat more qualitative results reported by Haszeldine.F The tetrahydrofuran-cyanogen iodide complex was studied only a t 27", using the same technique as for the the others. The results are given in Table I. The value of AHo for the formation of the complex is obtained from the temperature dependence of K by plott,ing R In K vs. 1/T, as shown in Figure 3. The resulting values of AHo and ASoZas(based on a standard state of 1 M concentrations) are listed in Table 11, together with values from other 1, complexes for comparison.
O.OI
-2.0
-
( V T ) x Id 3h0 340
360
360
Figure 3. The relationship between R In K and 1/T. Straight lines 1 and 2 correspond to EtOH.ICN and Et*O.ICN, respectively.
Discussion On the basis of the results shown in Table 11, we conclude that ICN and Iz have similar acceptor strengths. The enthalpies for the ICN complexes with alcohol and ether and AGO for the T H F complex are perhaps slightly more negative than for the corresponding I2complexes, as might be expected from the additional energy contribution due to the dipoleinduced-dipole term for the polar I C s complexes ~ I C N'v 3.76 in benzene solution3). However, the best interpretation of the results in Table I1 is apparently that I, and ICX possess almost identical acceptor strengths. It is of interest to note there that the points for AHo and ASo for the ICN complexes (after conversion to the same standard state) fit quite well on the line (23) The isosbestic point shifts at higher temperatures for the
CZHSOH-ICN solutions. When these solutions are cooled again to 2 8 O , the spectrum is identical with that observed at the start of the
experiment, indicating that the higher temperature changes responsible for the loss of isosbestic point are reversible. We do not understand this observation, but we have obtained the thermodynamic properties reported here from studies below 4 5 O , where the isosbestic point in Figure 2 remains constant.
Volume 71, Number 10
September 1967
HARID.
3292
Table I1 : Thermodynamic Properties for Some Complexes of ICN and of Some n Donors (Based on 1 M Solutions as the Standard State)
with
kcal/mole
-AP, cal/deg mole
0.79 0.74 1.54 0.11 2.34
6 . 5 -f 0 . 5 4.5 f0 . 5
19.1 12.7
... ...
...
... ... ...
-0.08 -0.09 0.65 0.11 3.37
4.2 4.6 =k 0 . 4 5.3 3.5 8.0
14.4 15.6 15.9 11.4 15.5
Kaw,
-AGO,
Acceptor
Donor
Solvent
l./mole
kcal/mole
ICN" ICN" ICN"
Et20 EtOH THF Dioxane Pyridine
%-Heptane n-Heptane n-Heptane CHCla C6Ho
3.8 3.5 13 1.2 51
Et20 EtOH THF Dioxane Pyridine
n-Heptane n-Heptane n-Heptane n-Heptane n-Heptane
0.87 0.86 3.0 1.2 290
ICN~ ICN~
LC Izd
LE 1 ~ 4
Le a
12
This research.
* From Popov, Humphrey, and Person.9
O
BISTAND WILLISB. PERSON
-AHo,
From M. Brandon, M. Tamres, and S. Searles, Jr., J . Am. Chem. Soc., From Julien.22 Recomputed from the summary by G. Brie-
82, 2129 (1960); M. Tamres and M. Brandon, ibid., 82, 2134 (1960).
gleb, ref 17, pp 124-125.
found empirically for 1, complexes24 when -AS" is plotted against -AH". We might also note here that the monomer-dimer equilibrium for ethanol in n-heptanez5 may complicate the interpretation both of the ICN and of the I2data for this donor. The concentration ranges of ethanol used to obtain K in both studies are roughly comparable, however, so that the comparison between I:! and ICN in Table I1 is probably relatively valid. It is of some interest to investigate quantitatively the changes in the 247-mp ICN band on complexing and to compare them with those for the corresponding changes in the visible band of 12. The intensity of the absorption of this band in the complex reported in Table I may be too high, since the increased overlap with the stronger band around 200 mp in the complex makes it difficult to obtain an accurate intensity for the 247-mp band above. However, examination of Figures 1 and 2 suggests that the overlap is not too great on the Hence, we have estilong-wavelength side of A,,. mated the f number and transition dipole, D,by the approximate formulasz6
f = 4.32 X 1 0 - g [ ~ m a x A ~ ~ / , ]
D = 0.0958[ e m a x A ~ i / , / v m a ~ ] (5) Here, we have estimated that An/*is two times the difference in reciprocal centimeters from the half-intensity point on the right-hand side of the band to the wavenumber of maximum absorption, vmrx. This estimate may be in error by about 5% or so, owing to the asymmetry of the band. However, the overlap by the absorption around 200 mp prevents us from measuring Av1l2 directly. The final results are given in Table 111. The Journal of Physical Chemistry
We see there that the "blue shift" and intensification of this band in complexes of ICN are indeed quite similar to those observed for the complexes of 12. The band for free ICN is intrinsically much weaker than the corresponding band in free I2and may be more susceptible to "inert solvent" effects. Hence, the ratio of intensities for ICN (complex/vapor) is greater than for 1 2 , but we see that D increases by about 0.2-0.3 debyelh for both I2 and ICN complexes. Finally, we may examine the reason for the failure to observe the charge-transfer band. We believe that the intensification observed for the complexed ICN in the region around 200 mp (Figures 1 and 2) is probably not great enough for it to be due to the chargetransfer absorption. Hence, ACT is probably less than 190 mp. If so, the charge-transfer frequency must come more than 12,000 cm-1 higher (or ~ V C Tis more than 1.5 ev greater) than the corresponding transition for I2 complexes. Such a large difference from I 2 is not obviously to be expected. The charge-transfer frequency is determined by:z6-z8 (1) the ionization potential of the donor, which is the same for ether complexes with I 2 and with ICN, for example; (2) GI, the stabilization of the dative-state energy, W1; GI should be roughly the same for I2 and for ICN complexes, assuming the configuration in the latter to be D- - -1(24) M. Tamres and M. Brandon, J . Am. Chem. SOC.,82, 2134 (1960). (25) U.Liddel and E. D. Becker, Spectrochim. Acta, 10, 70 (1957); E.D.Becker, ibid., 17, 436 (1961). (26) See ref 17; also cf. R.8.Mulliken and W. B. Person, Ann. Rev. Phy8. Chem., 13, 107 (1962). (27) R. S. Mulliken, J . Am. Chem. Soc., 74, 811 (1952). (28) W.B. Person, J . Chem. Phys., 38, 109 (1963).
CYANOGEN IODIDE COMPLEXES WITH SOMEn DONORS
3293
Table I11 : Changes in the Frequency and Intensity of the 247-mp Band and of the Visible Iz Band on Complexing
Amax,
mp
Vapor n-Heptane EtOH (CC14) Et10 (n-heptane) T H F (n-heptane)
520 520 443 462 4%
--
-ICN*
----Iodine’ Avl/a cm-1
3300 3200 4500 4100 (4500)8
%ax
832 918 1082 950 950
-
f
D
cm-1
mp
A& cm-1
0.012 0.013 0.021 0.017 0.018
1.15 1.19 1.41 1.29 1.40
...
249 241 217 220 225
8400 f 100 8200=!~200 9000 9000 9000
(vc
v d , C Xmax.
0 3300 2300 2800
(vc
-
klSX
f
D
cm-1
72 160 268 263 250
0.0026 0.0056 0.010 0.010 0.010
0.35 0.56 0.71 0.73 0.68
1300 5900 5300 4200
...
a From H. Tsubomura and R. P. Lang, J. Am. Chem. Soc., 83,2085 (1961), except for Iz vapor, which is estimated from E. Rabinowitch and W. C. Wood, Trans. Faraday SOC.,32,540 (1936); see also E. A. Ogryzlo and G. E. Thomas, J. Mol. Spectry., 17, 198 (1965); and for the T H F complex which is estimated from Brandon, Tamres, and Searles (footnote c of Table 11). This research. All solutions were in n-heptane. ’ Y O = wavenumber of maximum absorption in the vapor; Y C = wavenumber of maximum absorption in the solution, or of the complex in solution. This difference is probably accurate to about &500 cm-1, with the larger uncertainty applying A v i j 2 was measured for ICN vapor and in n-heptane solution; for the comin particular to the overlapping band in the complexes. ~ described in the text. The uncertainty in these plexes, where the overlap with the 200-mp band was a problem, we estimated A Y I / as Estimated. estimated values is believed to be f500 cm-l.
’
CN; (3) the resonance energies, XI and X o (=pZi/W1 W o ) ;(4) the electron affinity of the acceptor; and ( 5 ) Gothe stabilization (or repulsion) energy of the no-bond state, Wo. The latter three terms must contain the clue to the large difference in charge-transfer frequencies of ICN complexes, compared to the Izcomplexes. The electron affinity of C N is about 3.2 ev;2e the dissociation energy of the I-C bond is expected to be about 2.5 ev,30and the IC bond length about 2.00 A.30 Following the reasoning given for the halogensJZ8we predict the dissociation energy for the IC bond in (1CN)- to form I CN- to be about 1.25 ev, with T I C about 2.3 A and V I C about 240 cm-l. With these estimated parameters, the vertical electron affinity of ICN is predicted by the methods of ref 28 to be about 0.9 ev, or about 0.8 ev less than for 1 2 . In order to account for the remaining difference observed in the charge-transfer energies, consider Go. For I2 complexes with ethers, we should expect Goto be about +0.5-1.0 ev, by analogy with the values for amine-iodine complexes,28together with the knowledge that the 0-1 distance in ether-iodine complexes is relatively If the 0-1 distance in ethercyanogen iodide complexes is similar, we may still expect Go to be more negative than for the I2 complexes because of the increased electrostatic attraction between the polar ICN and the donor. This effect may be partially canceled, however, by the decreased resonance energy for ICN compared to I2 resulting from the in-
+
creased separation between W1 and Wo (see ref 28). However, this latter effect may also result in a longer 0-1 distance in the complexes with ICN, so that Go may be reduced essentially to zero. Hence, it would seem reasonable that the lower vertical electron affinity of ICN compared to Iz (0.9 ev, instead of 1.7 ev), together with a more negative Go (0 instead of +0.5 ev) resulting from the increased electrostatic attractive energy in the ICN complexes, may indeed account successfully for a shift of the chargetransfer frequency in ICN complexes toward higher energies by 1-2 ev, when compared to Izcomplexes.
Acknowledgments. Financial support from Public Health Service Research Grants No. GM-10168 and GM-14648 from the Division of General Medicine, PHS, is gratefully acknowledged. We are grateful to Dr. Larry Julien for the program for the Liptay calculation and for his cooperation elsewhere. It is a pleasure to acknowledge assistance from the University of Iowa Computer Center. (29) J. T. Herren and V. H. Dibeler, (1960).
J. Am. Chem. SOC.,8 2 , 1555
(30) (a) L. Pauling, “The Nature of the Chemical Bond,” 3rd ed, Cornel1 University Press. Ithaca, N.Y.,1960; (b) L. E.Button, Ed., “Tables of Interatomic Distances and Configuration in Molecules and Ions,” Special Publication No. 11, The Chemical Society, London, 1958. (31) 0.Hassel and C. Romming, Quart. Rev. (London), 16, 1 (1962).
Volume 71,Number 10 September 1087
