MILTONTAMRES AND JOHNM. GOODENOW
1982
Vapor Phase Charge-Transfer Complexes. I. Diethyl Sulfide-Iodine’
by Milton Tamres and John M. Goodenow Chemistry Depart&,
University of Michigan, Ann Arbor, Michigan
(Received May $3, 1966)
The diethyl sulfide-iodine complex has been studied in the vapor phase by a spectrophotometric method in the temperature range 94-127”. Several of the thermodynamic and spectral properties of the complex in the vapor phase differ significantly from those previously found in n-heptane solution. However, the differences for this case, which is among the stronger charge-transfer complexes with iodine, are not as pronounced as those recently reported for the weaker complexes of iodine with benzene and with diethyl ether. These general trends are consistent with hypotheses on solvent influence.
Introduction Spectral properties of charge-transfer (CT) or donoracceptor complexes and their theoretical interpretations have received wide attention in the past 15 years.2 The vast amount of spectral and thermodynamic data which has been accumulated is the result of studies made almost entirely on complexes in solution. The theory, on the other hand, uses properties which apply to the isolated species, a condition which is approached only in the vapor slate at low pressures. That spectral and thermodynamic properties are affected by the solvent, even those commonly regarded as “inert,” is apparent even in an early experimental report on the subject.* More detailed analyses of solvent effects confirm how large the influence of solvent can be.4 Very recently, several studies of CT complexes in the vapor phase have been reported: iodine complexes with benzene and with diethyl ether,6 carbonyl cyanide complexes with ethers and with aromatic hydrocarbons,6 and the tetracyanoethylene complex with p-xylene.’ These, generally, can be classified among the weaker CT complexes. The results make clear that important differences exist with regard to band position, equilibrium constant, and extinction coefficient. According to t h e ~ r y these , ~ ~ ~differences should be smaller for the stronger CT complexes. The strongest complex studied to date, diethyl sulfideiodine, was mentioned in a preliminary report.10 A more detailed account is given in this paper. Experimental Section Reagents. The source and purification of iodine,” The Journal of Physical Chemistry
of n-heptane,“ and of diethyl sulfide12have been given elsewhere. Vapor phase chromatography showed the absence of impurities in the diethyl sulfide and in the n-heptane. The iodine and the diethyl sulfide were kept in separate storage bulbs in a Pyrex vacuum line and prior to use the vapors of both were dried over phosphorus pentoxide. Apparatus and Procedure. The metering out and transfer of reagents into a cylindrical quartz cell were made using the Pyrex vacuum line. Teflon-threaded (1) Taken in part from the Ph. D. Thesis of J. M. Goodenow, University of Michigan, Dec 1965. This work was supported by a grant from the National Science Foundation, NSF GP-3691 Research, and also by a grant from the Horace H. Rackham School of Graduate Studies at the University of Michigan. (2) (a) G. Briegleb, “Elektronen-Donator-Acceptor-Komplexe,” Springer-Verlag, Berlin, 1961; (b) R. S.Mulliken and W. B. Person, Ann. Rev. Phya. Chem., 13, 107 (1962); (c) L. J. Andrews and R. M Keefer, “Molecular Complexes in Organic Chemistry,” Holden-Day, San Francisco, Calif., 1964. (3) H.A. Benesi and J. H. Hildebrand, J. Am. Chem. SOC.,71, 2703 (1949). (4) C. C. Thompson, Jr., and P. A. D. deMaine, J . Phya. Chem., 69, 2766 (1965). (5) (a) F. T.Lang and R. L. Strong. J. Am. Chem. SOC.,87, 2345 (1965); (b) F. T. Lang, Ph. D. Dissertation, Rensselaer Polytechnic Institute, Troy, N. Y., 1964. (6) (a) J. Prochorow, J. Chem. Phys., 43, 3394 (1965); (b) J. Prochorow and A. Tramer, ibid., 44, 4545 (1966). (7) M.Kroll and M. L. Ginter, J .Phys. Chem., 69, 3671 (1965). (8) N.S.Bayliss and E. G. McRae, ibid., 58, 1002 (1954). (9) 8. Carter, J. N. Murrell, and E. J. Rosch, J. Chem. SOC.,2048 (1965). (10) J. M. Goodenow and M. Tamres, J . Chem. Phys., 43, 3393 (1965). (11) M. Brandon, M. Tamres, and S.Searles, J.Am. Chem. Soc., 8 2 , 2129 (1960). (12) M. Tamres and S. Searles, J . Phys. Chem., 66, 1099 (1962).
VAPORPHASE CHARGE-TRANSFER COMPLEXES
valves (Fischer-Porter Co.) were used to avoid working with stopcock grease which adsorbs the reagents. The techniques for obtaining known concentrations of diethyl sulfide in the 10-cm absorption cell (2-cm i.d.) were similar to those used by Lang and S t r ~ n gLe., ,~ (1) determining the pressure and temperature of a known volume of vapor and (2) measuring out a t a fixed temperature a known height of liquid diethyl sulfide from calibrated capillary tubing. A check of one method against the other gave very good agree ment (within -1%). A further check was made by measuring out into the absorption cell a known amount of diethyl sulfide by the capillary method, adding a known volume of n-heptane, and determining the sulfide concentration from the known spectrum of diethyl sulfide in n-heptane solution. Again, agreement was good. The concentration of iodine was measured approximately by allowing the iodine to sublime a t room temperature into a bulb of known volume. The iodine was condensed in the absorption cell together with the sulfide a t liquid nitrogen temperature and the cell was sealed off from the vacuum line. After the vapor phase data had been taken, a more precise determination of the iodine concentration was made. This was done by "cracking" open the cell and adding a known volume of n-heptane. Since thermodynamic values for the diethyl sulfideiodine complex in nheptane were known from earlier work,12 the solution data were used to calculate the initial concentration of iodine in the vapor phase. The diethyl sulfide-iodine adduct is not very volatile and requires elevated temperatures to get appreciable concentrations in the vapor phase. This was accomplished by designing a twin aluminum jacket which held the cell with complex in one compartment and a reference cell in the other. For complete insulation, aluminum covers were placed over the side arms of the cells and quartz covers were inserted in grooves cut in each end of the thermojacket to protect the cell windows. The thermojacket was wrapped with electrical heating tape and the temperature was controlled by powerstats. Temperature equilibrium could be established in about 20 min. Absorbancy readings over the wavelength range were made in about 0.5 hr, during which time the maximum temperature fluctuation was of the order of 1". Temperatures were measured with a thermocouple which could be inserted into the side arm of the reference cell and moved to various positions within the cell. A special aluminum and Bakelite housing was constructed for the cells and thermojacket assembly. To minimize reflection of stray light, the thermojacket
1983
and aluminum parts of the housing were anodized and dyed black. Thermospacers were inserted next to the body of the spectrophotometer (Beckman DU) and next to the phototube housing for cooling. The Concentration of diethyl sulfide varied from 2.8 X 10-3 to 1.6 X M and that of iodine from 1.4 X 10-5 to 5.8 X M . Spectra were recorded a t 250-500 mp over a temperature range of 94-127", with the absorbancy readings a t A,, (290 mp) ranging from a high of nearly 0.7 to a low of nearly 0.1. A t the end of a run, return was made to the initial temperature, and readings at Amax were taken as a check on the reproducibility of the data. Calculations. The spectrophotometric equation which applies to this study id3 CdCz - CB --
A - Ao
E,
+ Cz +
-
€18
1 Kc(~c-
€12)
A - A0 b(ec - EIJ' (1)
where CB and CZ are the initial concentrations of sulfide and of iodine, respectively, b is the cell length, ec and eI, are the extinction coefficients of the complex and of free iodine, respectively, A is the absorbancy, A. is E I ~ ~ Cand Z , K, is the association constant. At A,, (290 mp), ern is quite small, so that the A. correction is very small. Calculations of Kc and e, were programed for the IBM 7090 computer. The error limits of K, were determined using the method of Fieller.'* For work in solution at room temperature, reducing eq 1 from quadratic to linear form by dropping the last term produced noticeable error in K,,12 but in the present work a t elevated temperature, a t which condition K, is smaller, dropping the last term made very little difference, in accordance with expectation. l 3
Resulk There are shown in Figure 1, for typical concentrations used in studying the ultraviolet region, the spectra of free diethyl sulfide and free iodine, together with the spectrum which results when the components are mixed. There is evident the intense CT band with a maximum a t 290 mp, a region which is relatively optically clear with respect to the components. The absorption of diethyl sulfide vapor becomes appreciable only at wavelengths less than 270 mp for the concentrations used in this study and no correction for free sulfide is necessary in reading the absorbancy a t Amax of the complex. The sulfide vapor showed a small increase in absorbancy with increase in temperature over the range 63-1 18" in the limited wavelength region (13) M. Tamres, J . Phy8. Chem., 6 5 , 654 (1961). (14) E. C . Fieller, Appendix Suppl. J . Row. Stat. SOC.,7, 1 (1940).
Volume 71, Number 7 June 1967
MILTONTAMRES AND JOHN M. GOODENOW
1984
Figure 1. Spectra in the vapor phase at 100' in a 10-cm cell of 1, diethyl sulfide = 2.78 X M; 2, iodine = 4.52 X lo-@M ; and 3, diethyl sulfide iodine (1 2).
+
+
5 ( [Et2S]
investigated (270-250 mp) where the sulfide begins to absorb. This is indicative of temperature broadening. In the region where the CT complex is found, there is a small absorption by free iodine, which results in a very small correction in determining the absorbancy of the complex alone at A., As shown in Figure 1, the iodine has a weak absorption band with a maximum below 290 mp. This band has been reported not to obey Beer's law and has been attributed15 to the equilibrium
212
-
A
14
(2)
At 290 mp, the main contribution of the iodine absorp tion is due to the species Is rather than 1 4 . The extinction coefficient of the former15J6at 290 mp is 10.2 1. mole-' cm-l, the value used for in eq 1. For the concentrations of iodine used in this work, A0 in eq 1 did not exceed 0.005 and averaged closer to 0.002. Data for seven donor-acceptor concentrations are plotted in Figure 2 using a modified Benesi-HildebrandScott equation, obtained by dropping the last term in eq 1. The best lines through the points were determined by the method of least squares. Good straight lines also were found using the Benesi-Hildebrand equation.g The points in the vapor plot show slightly more scatter than those obtained for this complex in n-heptane solution.12 A contributing reason for the greater scatter in the vapor study is due to the indirect method of determining the iodine concentration separately at the end of each vapor run by using the spectrophotometric method after the complex was dissolved in solution. This is less precise than the method used in solution work where the iodine concentration is determined directly by dilution of a common stock solution of known concentration. The Journal of Physical Chemistry
I5
IO
+ [I2] 1
X
IO3
Figure 2. Modiiied Benesi-HildebrandScott plot of vapor phase data for the diethyl sulfide-iodine complex in the vapor phase.
The values of K Oand e, determined from the absorbancy data at A,, are given in Table I. These results are based on a calculation of the initial iodine vapor concentration assuming that in n-heptane solution12 a t 25" K O = 180 1. mole-' and e, = 27,600 1. mole-' cm-l, these values being in a range similar to those reported by several investigators. l7 The solution data,12 however, in one case had led to values at 25" of K , = 200 1. mole-' and e, = 24,800 1. mole-' cm-' which, of course, would alter the vapor phase results somewhat. To determine the magnitude of the effect, a second set of calculations based on the latter values was made and is given in parentheses in Table I. It is seen that a 10% change in K Oand e, in the solution data produces a larger percentage change in the vapor phase results, but the general order of magnitude remains similar. Calculations involving eq 1 to evaluate K Oand eo are based on the assumption that only a 1 : 1 complex exists. In the present study, this assumption is supported qualitatively by the narrow (half-width -5700 cm-l) and nearly symmetrical shape of the diethyl sulfideiodine CT band, published previously.l o But this feature alone is not conclusive. Even the fact that the (15) G. Kortum and G. Friedheim, 2. Naturforsch., 2a, 20 (1947). (16) M. Tamres and J. M. Goodenow, unpublished work. (17) (a) H. Tsubomura and R. P. Lang, J . Am. C h . SOC.,83,2086 (1961); (b) M. Good, A. Major, J. Nag-Chaudhuri, and 8. P. McGlynn, ibid., 83, 4329 (1961). Calculations here were based on the Benesi-Hildebrand equation, which for this strong interaction would result in small error.**
VAPORPHASE CHARGE-TRANSFER COMPLEXES
1985
~~
~
~~
~
Table I: KOand eo for Diethyl Sulfide-Iodine in the Vapor Phase 7
Temp, O C
At Xu.
Liptay analyais
280 mp
Ko,'llb
KO
t0,41b
1. mole-'
1. mole-1 em-1
1. mole-1
eo at 280 z a p , 1. mole-1 om-1
94
17.8 f 6.0 (22.9f 6.7)
12,500f 4,000 (9,480f 3,100)
17.6 f 6.5 (22.9f 7.2)
12,600 (9,440)
100
19.4 f 2.5 (24.0f 2.6)
10,000f 1,000 (7,910f 830)
19.5 f 2.7 (24.0f 2.7)
10,000 (7,930)
105
13.7f 2.5 (17.8f 2.6)
11,300f 1,600 (8,460f 1,200)
C
C
C
C
115
10.2f 1.8 (14.1f 1.8)
12,300f 2,200 (8,640f 1,100)
9.6 f 1.8 (13.3f 1.7)
122
4.7 f 2.8 (8.8f 2.8)
21,200 f 12,600 (11,200f 3,200)
C
C
C
C
127
9.9f 1.2 (13.3f 1.3)
9,600f (6,930f Av
760 430)
11,200df 1,900 (8,770f 1,640)
13,000 (9,080)
9.3 i 1.3 (12.2f 1.7)
10,300 (7,500) Av
11,480 f 1,330 (8,490f 770)
'
a Error limits are for the 50% confidence interval. The values not in parentheses were obtained by basing the calculation of initial iodine concentration on the solution values in n-heptane at 25" of KO = 180 1. mole-' and eo = 27,6001. mole-' cm-1; those in parenReadings taken at A, only. theses are based on the solution values of KO = 200 1. mole-' and eo = 24,800 1. mole-' cm-I. Average does not include results at 122".
Benesi-Hildebrand and Benesi-Hildebrand-Scott type of plots at the single wavelength Amax showed good linearity, whereas plots based on 2 : l complex only and on 1 :2 complex only (less likely for the concentrations used) did not, does not eliminate completely the possibility of having a second complex of 2: 1 ratio present in smaller concentration. 18,19 The second complex would produce a systematic variation in equilibrium constant with wavelength and it has been suggested that this be used as a criterion to determine its presence.'* To test this possibility, a LiptayZ0 analysis of the data a t eight wavelengths over the range 280-310 mp was made. The constancy in the columns of the D m k matrix indicates the dominant presence of the 1 : l complex and only small, random variation in K , was obtained over the wavelength range. The L i p tay average value for K Oand the value for e, are given in Table I and are seen to fall within the error limits No Liptay analysis for the values obtained a t Xm,. was made at 105 and 122" because data for these two Other temperatures were recorded only a t A.,, work which shows the predominant formation of a strong 1 :1 CT complex with iodine is that of Klaboe, et aL121who used the Job method of continuous variations to study the interaction with triphenylarsine in carbon tetrachloride. It is well known that the product Koeocan be de-
termined more precisely than either term alone, any error in one producing a comparable error in the other. The wide variation in e,, which appears to be random with temperature, results in a nonsystematic temperature dependence of KO. For such cases it is best to assume that ec is independent of temperature and to plot log Koecus. 1/T to calculate the heat of reaction. An equivalent procedure22is to plot [log K e f o- log eOav] us. 1/T. Such a plot is shown in Figure 3, where the term in brackets has been abbreviated as log Keav. The calculated thermodynamic values are given in Table I1 together with the spectral characteristics, and comparison is made to the results found in n-hep tane solution. The values in parentheses are based on the set of data given in parentheses in Table I. It is apparent that any uncertainty in the selection of solution values for K Oand e, to calculate iodine vapor con-
(18) G. D. Johnson and R. E. Bowen, J. Am. Chem. SOC.,87, 1655 (1965) (19) C. C. Thompson, Jr., Abstracts, 151st National Meeting of the American Chemical Society, Pittsburgh. Pa., March 1966,paper 112N. (20) W. Liptay, 2. EZektrochem., 65, 375 (1961) : ala0 Chapter 12 of ref 2a. (21) E. Augdahl, J. Grundnes, and P. Klaboe, I m o . C h m . , 4, 1475 (1965). (22) M. Tamres, J . Phys. Chem., 68, 2621 (1964). I
Volume 71, Number 7 Juna 1067
MILTONTmms AND JOHNM. GOODENOW
1986
Table II: Vapor and Solution Spectral and Thermodynamic' Properties of Diethyl Sulfide-Iodine
I .4
n-Heptane solutionC
Vaporb
1.2
Xwx, mic eCav, a ,t A
s
303 26,400
K,,,,, 1. mole-'
16.5d (20.7)
9.38;" 12.5'
K4*,1. mole-'
226' (285)"
200; 180U
2.82 (2.98)
1.66; 1.75'
3 * 75' (3.89)'
3.14; 3.08O
cm-I
Y
3
290 11,200 (8770)
1. mole-'
1.0
0.8
r"
-AGoa,a, kcal mole-'
I
I
2.5
2.6
IO' T
I 2.7
AGO^@, kcal mole-'
1
2.8
-ASo, eu
Figure 3. Log KO€,- log ecBv (= log Kea,) vs. 1/T for the diethyl sulfide-iodine complex in the vapor phase.
centration, while affecting AGO slightly) does not affect the other thermodynamic quantities.
Discussion The CT complex in its ground (N) and excited (V) states can be described approximately by the wave functions. = a%
+
(3)
and q v = aW1 - b*qo
(4)
where *O and \YI are the "no-bond" and "dative bond" wave functions) respectively.28 The transition energy h v C T for the process !PN -.*. 9~ in the case of weak complexes has been approximated as2b.24
(5)
and, for strong complexes, as 2b,26
Here
1 5 . 0 3 ~1.1' (14.8 f 1.0)'
19.4 f 2.0' 17.6 f 0.5"'
-AHo, kcal mole-'
8 . 4 f0 . 4 h ( 8 . 4 f 0.4)'
8 . 9 f 0.6' 8 . 3 f0 . 2 g 3 h
kcal mole-'
7 . 7 f 0.4' (7.7 f 0.4)'
8 . 9 f 0.6' 8 . 3 f 0.2O*'
-Ah'',
Relative to a standard state where the concentrations of donor, acceptor, and complex are 1 M. Where two values are given, one without and one with parentheses, that without parentheses was obtained by basing the calculation of initial iodine concentration on the solution values in n-heptane at 25' of K , = 180 1. mole-' and eo = 27,600 1. mole-' cm-1 and hhat with parentheses is based on the solution values of K , = 200 1. mole-' and e, = 24,800 1. mole-' cm-' (see Table I). Reference 12. Read from the line in Figure 3, as determined by the method of least squares. Extrapolated from ultraviolet data assuming constant AE". Extrapolated from study of blueshifted iodine band in visible region. ' From data in visible region. Standard error.
'
'
donor; EA^ is the vertical electron a n i t y of the acceptor; E,, the coulombic energy, is equal to -e2/r where e is the electrostatic charge and r is the intermolecular distance; Sol is the overlap integral; W1 and Wo are the energies associated with Q1 and \ko, respectively; and PO and PI are terms which contain overlap and exchange integrals.26 An approximation which has been utilized to estimate the polarity of weak CT complexes isz7 ~~
ci =
EA^
- E , + wa
cz = Po2
+ Pi2
A =
- Wo
W1
(7)
(8) (9)
where ID"is the vertical ionization potential of the The Journal of Physical Chemistry
(23) R. S. Mulliken, J . Am. Chem. Sac., 74, 811 (1952). (24) 8. H.Hastings, J. L. Franklin, J. C. Schiller, and F. A. Matsen, {bid., 75, 2900 (1953). (25) H.Yada, J. Tanaka, and 9. Nagakura, Bull. Chem. Sac. Japan, 33, 1660 (1960). (26) R. S. Mulliken, J . Am. Chem. Sac., 74, 811 (1952). (27) J. A. A. Ketelaar, J . Phy8. Radium, 15, 197 (1954).
VAPORPHASE CHARGE-TRANSFER COMPLEXES
1987
tion are comparable. Some difference has been reported for ben~ene-iodine~and quite a large difference has been found for p~ylene-tetracyanoethylene.~-~O where AH" is the heat of reaction for complex formaAny comparison of AH" should take into account the tion, v,, is the fLequencyat the CT band maximum, fact that solution and vapor work have been carried and a and b are the coefficientsin eq 3. out in different temperature ranges and there may The parameters in eq 5 have been determined for be a small temperature dependence of AHo over this iodine complexes with a wide variety of weak donors2a*b temperature range. For the donor and acceptor sepaand in eq 6 for iodine complexes with amines.2bg26 rately, AH"(so1v) can be determined from the heat of Reasonable results for the polarities have been obtained solution and the heat of vaporization (or heat of subliusing eq 10 for iodine complexes with aromatic hydromation in the case of iodine). The sum of these enercarbons2' and with ethers.28 In all instances these gies is quite large, being several times that of the heat equations have been applied to studies in solution in of reaction. The heat of solvation of the vapor CT which a similar of type solvents (predominantly complex must also be large and, although not determinn-heptane) was used, where it might be expected that able directly, especially for weak complexes, can be solvation factors would parallel other trends in the calculated if all other terms in eq 12 are known. Thus, donor-acceptor properties. It is evident that the the heat of solvation of benzene-iodine vapor in properties of the donor, acceptor, and complex will be carbon tetrachloride is of the order of - 15 kcal/mole.31 modified by solvent interaction, thereby influencing the The Born-Haber-Fajans cycle for the charge-transindividual terms in eq 5-10. Thus, the parameters fer energy is in eq 5 and 6 and also the polarities of the complexes will be different in the vapor than in solution and, in fact, will be different for each solvent. aolv A E v solv AEN (13) Vapor and solution thermodynamic and spectral h vCT (soh) DZN(so1n) ----+- DZv(so1n) values are related by the Born-Haber-Fajans cycle. When donor (D) and acceptor (Z) form a complex (DZ), which gives any thermodynamic function (AX') will differ in the two phases as shown hvcT(g) = hVc'r(Soln) [ a B - G N l (14)
.1
D(soln)
+
Z(so1n)
from which AXo(g)
=
A X 0 (soh)
DZ(soln)
AXO(so1n) -
[AXO~z(soiv) - AXO~(s0iv)- AXoz(Bolv)I (12)
If the thermodynamic function is to be the same in both phases, the term in brackets would have to be zero; but in general, the sum of the changes due to solvation in donor and acceptor separately does not cancel that of the change in the complex. .This undoubtedly is the case for AGO where the recent vapor phase studies of various donors (aromatic hydrocarbons and ethers) with several acceptors (iodine, tetracyanoethylene, and carbonyl cyanide), as well as the present study on diethyl sulfide-iodine, consistently give equilibrium association constants which at comparable temperatures are appreciably higher in the vapor phase, even by as much as a factor of 30.29 I n the present study and in the case of diethyl etheri ~ d i n eAHo , ~ in the vapor phase and in n-heptane solu-
.1
In all of the cases where spectral data on CT complexes have been obtained in both phases, there has been observed a red shift in going from vapor to solution. B a y l i ~ shas ~ ~discussed the effect of solvent in producing a polarization red shift. This effect is related to the polarizability of the solvent and, indeed, a correlation has been noted of CT band position of aromatic hydrocarbon-TCNE complexes with the refractive index of the solvent.33 Further, for these complexes, it has recently been shown for similar(28) M. Tamres and M. Brandon, J . Am. Chem. SOC.,82, 2134 (1960). (29) There is always a problem in determining a precise K , because of the dependence on evaluation first of 6,. Where the scatter of data is greater in a Benesi-Hildebrand plot, as is the case for some of the vapor phase work compared with that in solution, it is even more difficult to evaluate K,. Nevertheless, the difference between vapor and solution Kc values seems quite definite. (30) Large differences also exist for other polymethyl-substituted benzene-TCNE complexes: M.Kroll, private communication. (31) Calculation is based on ref 5; F. R. Bichowsky and F. D. Rossini "Thermochemistry of Chemical Substances," Reinhold Publishing Gorp., New York, N. Y., 1936: K. Hartley and H. A. Skinner, Trans. Faraday SOC.,46, 621 (1950); M.Tamres, J . A m . Chem. SOC., 74, 3375 (1952). (32) N. S. Bayliss, J . Chem. Phus., 18,292 (1950). (33) H. M.Rosenberg and D. Hale, J. Phys. Chem., 69, 2490 (1965).
Volume 71, Number 7 June 1967
MILTONTAMRES AND JOHN M. GOODENOW
1988
type solvents that there is a linear correlation of the CT energy with the refractive index function (n2- 1)/(2n2 l), the extrapolation of which to n = 0 corresponds to the CT energy found in the vapor
+
The interaction of solvent is different for the ground and excited states of solutes and depends on the relative polarities of these states.8 When the dipole moment of the excited state is larger than that of the ground state, there will be greater solvent stabilization of the former relative to the latter, which should lead to a larger spectral red shift in going from vapor phase to solution. Since the difference in dipole moment between the excited and ground states is larger for weak CT complexes than for strong, there should be a decreasing shift AT or, better, a decreasing relative shift AT/T in going from weak to strong complexes. For the available data on iodine complexes presented in Table I11 this seems to be the case,a5but the data are much too few to allow generalization. Apparently this does not hold for carbonyl cyanide complexes.s When the polarities of the ground and excited states are nearly alike, then M N'v MVand there would be little shift in ~ V C Tbetween vapor and solution. Further, the potential energy curves of the ground and excited states would have comparable shapes and minima, thereby giving CT band shapes which are nearly symmetrical. Comparison of the relative shapes
Table 111: Comparison of e,,, for Iodine Complexes in the Vapor Phase and in Solution Fg,
Donor
om -1
aoln, cm-1
A?, cm-1
A?&
Benzene
37,300"
33,600' 34,800'
3700 2500
0.097 0.067
Diethyl ether
42,800"
39,700d 40,300"
3100 2500
0.074 0.057
Diethyl sulfide
34,500
33,000'
1500
0.044
Reference 5. * I n n-heptane, ref 3. I n benzene, ref 24. I n n-heptane, ref 28. " I n diethyl ether, ref 15. 'In n-heptane, ref 12.
of the diethyl sulfide-iodine CT band in the vapor phase and in n-heptane solutionloshows them to be practically superimposable except for the shift in band maximum. They are only slightly asymmetrical, with the slope on the low-energy side decreasing a little more slowly than On the high-energy side' The vapor phase CT band for iodine with diethyl ether,5 still more asymThe J o u m l of Physical Chemistry
metrical than that with diethyl sulfide, also decreases more slowly in slope on the low-energy side. A similar finding for triphenylarsine-iodine in carbon tetrachloride has been noted.S6 These are among the n --t u* complexes. From accumulated data in solution for predominantly T+T* complexes, Briegleb and Czekallaa7note that the CT band is asymmetric with the slope decreasing more slowly on the highenergy side. The absolute intensity of the CT band in the vapor and in solution is considerably different. In all cases, ec is smaller in the vapor differing from the solution value by a factor of between 2 and 3 for diethyl sulfide-iodine, diethyl ether-i~dine,~and pxylene-tetracyanoethylene,? and by a factor of about 10 for ben~ene-iodine~ and CO(CN)2with diethyl ether and with benzene.6 Since K , is found to be larger in the vapor phase, the product Koe, is closer in proportion in the two phases than is either term ~ e p a r a t e l y . ~ ~ Lang and Strong have pointed out that these differences in the two phases can be accounted for in two ways. First, they say there is the possibility that the large extinction coefficient in solution is mainly due to contact charge transfer, especially for benzeneiodine. I n a solvent, a "cage effect" is responsible for the large concentration of contact charge transfer and this effect is negligible at reasonably low pressures in the vapor phase. Second, as proposed recently by Murrell, et U Z . , ~ in solution the solvation process competes with that of complexation. Omission of solvation leads to ec values which are too high and K , values which are too low. Application of solvent corrections to solution data for benzene-iodine leads to e, and K , results which are in qualitative agreement with those found in the vapor phase. The effect either of contact charge transfer or of solvation should be smaller for the stronger complexes. First, since contacts are random, the weaker complexes may have several different configurations for which contact charge transfer can occur. Stronger complexes favor a single configuration and, hence, there would be a smaller contribution from contact charge transfer. If for diethyl sulfide-iodine, it is assumed that there is very little contact charge transfer in the vapor phase and also very little in solution, it seems reasonable that the e, values in the two phases (34) E.M.Voigt, J . Phys. Chem., 7 0 , 598 (1966). (35) Possibly, this may not hold for iodine complexes with olefins; ref 39. (36) E.Augdahl, J. Grundnes, and P. Klaboe, I m r g . Chem., 4, 1476 (1965). (37) G. Briegleb and J. Czekalla, Z . Physik. Chem. (Frankfurt), 24, 37 (1960).
VAPORPHASE CHARGE-TRANSFER COMPLEXES
for this complex would be in closer agreement than, say, for the benzene-iodine case. Second, Murrell's equationsg for K O and e, when there is solvation show that the solvent should not have as much effect OK strong complexes as it does on weak ones. Thus, there should be closer correlation of K , and e,, in agreement with experiment. Another possibility for the difference of e, in the two phases has been proposed by Prochorow and Tramer as being6b due to the internal pressure of the solvent which would decrease the intermolecular distance between donor and acceptor in the complex, thereby increasing the overlap integral and, hence, e,. It might be supposed that compression by solvent would raise the energy of the complex from its potential energy minimum, thereby decreasing AH". Perhaps this is more pronounced for the T-T* complexes, such as p xylene-tetracyanoethylene, which may be more compressible than the n-a* complexes, but the absence of any difference in AH" between vapor and solution for the latter type of complexes makes it doubtful that compressibility is the dominant factor. An increase in overlap due to compression of the complex should increase CZ (eq 8) and, as seen in eq 5 for weak complexes, should produce a blue shift. Since this is opposite to what is observed in going from vapor to solution, there must be a change in some of the other terms as well. A red shift would occur if the sum -C1 C,/(ID' - C1) becomes more negative (assuming ID^ is unaffected in the two phases). This decrease could occur by increasing Claswhich, according to eq 7, could arise from compression by decreasing the coulombic energy term E,, but it is likely that the solvent affects the other terms in eq 5-10, as mentioned earlier, and the net shift is a composite of all these changes. The visible band of diethyl sulfideiodine in n-heptane has been reported by others'' and was investigated in our laboratory as well.' There is found a pronounced blue shift with a maximum at about 437 mp and is attributed to the perturbed molecular iodine which results from complexation. Effort was made to study the visible band in the vapor phase. As seen in Figure 4, no shifted maximum in the visible region was observed. For the con-
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1989
aoo}
0
400
1
I
500
600
Xhpl
Figure 4. Apparent extinction coefficient of iodine ( E = A/b[I&) when the iodine is 28% complexed: 1, in the vapor phase at 127'; 2, in n-heptane solution at 25'.
centrations used, according to the data in Table I, the iodine should have been complexed to the extent of 28% a t 127". An equal extent of complexation in n-heptane solution at 25" definitely shows the shifted band. Similar observations have been made by Lang and Strong6 on iodine complexes with benzene and with diethyl ether and by Van Tongerloo and Backs9 on iodine complexes with olefins. It is not likely that the absence of a shift is due to experimental diffculty. Diethyl sulfide, unlike benzene and diethyl ether, does not undergo secondary reaction with iodine at elevated temperatures, as evidenced by the study of the complex in the ultraviolet region. Being the strongest donor of the three, it is the most favorable for studying the visible shift. Even if the reported value of K , in the vapor phase were too large by a factor of 2,40there should be sufficient iodine complexed to show the effect, if vapor and solution were similar in behavior. Perhaps, as attributed by Bayliss and c o - w o r k e r ~ , *the ~ ~pres~ ence of a solvent cage is responsible for the blue shift, which would be absent in the vapor. This is a point which requires further study.
z/z
(38) A minimum occurs in eq 5 when I g V - Ct = and, for weak complexes, CI < (39) A. Van Tongerloo and M. H. Back, private communication. (40) Thie is not very likely, since this would result in eo being about the aame in the vapor and in solution, (41) N. 8. Bayliss and A. L. G. Rees, J . Chem. Phys., 8, 377 (1940).
my.
Volum 71, Number 7 June 1987
