SPECTROSCOPIC STUDY OF SOMECBr4 COMPLEXES tive functions do not imply better agreement than can be expected from reference to errors reported for experimental observations. An interchange of optimum force constants does not significantly reduce the efficiency of prediction of the two thermodynamic parameters studied. Either of the optimum pairs for a given fluid predicts transport data relatively well, although this latter capability is poorer than with any of the local optimum pairs generated with transport objective functions.
2309 I n summary, Lennard-Jones force constants for methane and argon mere derivable from thermodynamic data by direct optimization techniques, without risk of multimodality. Optimum force constants are not a strong function of the choice of thermodynamic parameters, so that pairs generated with one type of data can be applied to the calculation of other quantities. Transport data lead to multimodal, insensitive objective functions and are not a suitable source for the optimization process.
A Spectroscopic Study of Some CBr4 Complexes by Donald A. Bahnick, William E. Bennett, and Willis B. Person Departments of Chemistry, University of Iowa, Iowa City, Iowa and University of Florida, Gainemille, Florida 38601 1 (Received November 81, 1968)
Results are presented from an investigation of the Raman and ultraviolet spectra of CBr4 in solutions containing benzene and from an ultraviolet investigation of the tetrahydrofuran-CBr4 system. The changes occurring in the ultraviolet spectrum for these systems have been analyzed quantitatively to determine the formation constants for one-to-one donor-CBrd complexes. Relative Raman intensity values are presented for three of the four CBr4 fundamentals in benzene-CBr4 solutions. These data were analyzed and interpreted in terms of complex formation, and the formation constants from this analysis were compared to those from the ultraviolet study. The two sets of formations constants agree, leading to the conclusion that formation of the same chemical species causes the observed changes in both the ultraviolet and the Raman spectra of CBr4 in benzene. The changes in the Raman intensities for CBr4 in solution with benzene are analyzed in terms of changes in the bond polarizability parameters, with results similar to those found earlier for solutions of Cc4 in donor solvents. The nature of the benzene-CBr4 interaction is discussed, and it is concluded that it is probably primarily electrostatic.
Introduction Evidence for the probable existence of complexes between CBr4 and certain electron donors has come from several kinds of study. For example, phase diagram studies have shown the existence in the solid state of one-to-one complexes of benzene-CBr4 and p-xyleneCBr4.2a From the X-ray structure determinationzb of the p-xylene-CBr4 complex, chains of alternating xylene and carbon tetrabromide molecules a.re found in the crystal, with bromine atoms from different CBr4 molecules centered, one on either side of the ring. Each CBr4 molecule is associated with two rings. These bromine atoms lie on perpendicular lines passing through the centers of the rings with the distance from the center of the ring to the bromine atom being 3.36 A. This distance is the same as that reported for the benzenebromine complexa and is somewhat less than the sum of van der Waals radii (3.65 A). A study of the nuclear quadrupole resonance (nqr) spectrum of this complex has been interpreted as indicating that there is little or
no transfer of charge from xylene to the bromine However, this interpretation has been questioned in a similar study by Gilson and O’KonskiJ4bwho suggest that the nqr results are also consistent with a small amount of charge transfer which is masked by the effects of weak “intercomplex” interactions in the solid. Additional evidence indicating CBr4 complex formation comes from the changes observed in the ultraviolet and infrared spectra of hexamethylbenzene in solutions containing CBr4.5 Tramer has shown that the ultraviolet absorption by CBr4 in solutions containing ben(1) Reprint requests should be addressed to Dr. Person at this address. (2) (a) A. F. Kapustinskii, Bull. Acad. Sei. USSR,435 (1947); (b) F. Strieter and D. Templeton, J . Chem. Phys., 37, 161 (1962). (3) 0.Hassel and K. Stromme, Acta Chem. Scand., 12, 1146, 1781 (1958). (4) (a) H. 0. Hooper, J . Chem. Phys., 41, 599 (1964); (b) D.F. R. Gilson and C. T. O’Konski, ibid., 48, 2767 (1968). (6) F. Doerr and G. Buttgereit, Be?. Bunsenges. Phys. Chem., 67,861 (1963). Volume 73, Number 7 J u l y 1969
D. A. BAHNICK, W. E. BENNETT, A N D W. B. PERSON
2310 sene, toluene, p-xylene, or 2-chloronaphthalene is enhanced and shifted from its absorption in hexaneU6 A linear relationship was found between the frequency of the absorption edge and the ionization potentials of the aromatic compounds, indicating weak electron donoracceptor complex formation between these aromatic donors and CBr4. Tramer estimated equilibrium constants for the formation of one-to-one complexes from these spectrae6 We have shown in investigation of the Raman intensities of the fundamental vibrations of cc14 in varicm electron-donor solvents that intensity changes occur for CC1, in the presence of certain aromatic and amine ~ o l v e n t s . ~Anticipating that CBr4 may interact more strongly with these donors than does CC14, we have investigated the benzene-CBra system. We have also attempted further studies of solutions containing piperidine (an amine donor) and CBr4. However, in the latter system an obvious reaction took place during spectral measurements, to form a white precipitate (probably the amine hydrobromide), so we confined our studies of the Raman spectrum to the benzene-CBrr system. To determine the formation constants of CBr4 complexes ultraviolet studies were made of the benzeneCBr4 complex and of the tetrahydrofuran (THF-CBr4) complex. The latter system was picked because it has been found that oxygen-containing donors interact more strongly with certain acceptors such as I, than does benzene.s The four fundamental vibration frequencies of CBr4 are listed in Table I. The 128-cm-’ CBr4 band appears on the shoulder of the intense mercury exciting line; consequently, it was not possible to make accurate intensity measurements on this band. However, the absolute Raman intensities of the other fundamentals were studied and interpreted in terms of bond polarizability derivative uarameters for the nlolecule, paralleling the treatment of the intensities of CClf and of the ICN comple~es.~ Table I : Fundamental Vibration Frequencies of CBr4 in Cyclohexane Solution Roman Description of modea
symmetryb
vl, C-Br stretch
a1
CBrz bend v ~ C-Br , stretch v,, CBrz bend
e f f
v2,
shifts? om-’
269 vs 128 s 673 ms 182 s
Qualitative (for quantitative description see normal coThis ordinates). b Classification according to T d point group. study; vs means “very strong,” etc.
Model 81 spectrophotometer with a mercury arc lamp (4348 A) as the source of excitation. Our procedure for the Raman intensity measurements has been described previ~usly.~In the present study the 804-cm-’ cyclohexane band served as an internal intensity standard. The Raman intensities were measured on ternary solutions containing benzene, CBr4and the cyclohexane internal standard, as described previously.7~9 Cyclohexane was purified by shaking several times with sulfuric acid, followed by distillation. Benzene was fractionally distilled, and T H F was purified by passing through a column containing a 3: 1 mixture of Celite and activated A1203. The CBr4 was Eastman Spectro Grade quality, a white solid which was used without further purification. However, upon exposure to strong ultraviolet light, CBrr slowly decomposes in solution to form bromine, among other products. Thus, we adopted the procedure used by Long and Milner’O to make Raman intensity studies of solutions of CBr4in CC14. Their procedure consisted of the addition of enough allyl alcohol (-1 mol %) to each solution to remove any bromine formed by decomposition during the course of recording the Raman spectra, so that the solutions remain colorless. I n our study the concentrations of allyl alcohol in the solutions varied from approximately 0.02 to 0.05 M , depending on the amount of CBr4 present. The CBr4 solutions were kept completely away from exposure to ultraviolet light as much as possible so that no CBr4 decomposition is believed to have occurred. However, if some decomposition did occur, then it is possible that we have overestimated CBr4 concentrations by as much as 27”. If the extent of decomposition were greater, then we would have observed the formation of colored solutions due to excess Brz. The ultraviolet spectra were recorded on a Beckman DK3A spectrometer as described elsewhere.8sQ We used a set of UV-01 short path length microcells (from Limit Research Carp.) using quartz windows separated by a Teflon spacer, with a path length of 0.1 mm. The path lengths were measured by observing interference fringes; they did not vary by more than *l%, even after completely disassembling them for cleaning. Benzene dissolved in cyclohexane absorbs to some extent throughout the 230-270-m~ region. It was necessary to cancel out this absorbance in our investigation by adjusting the concentration of benzene in our cyclohexane in the reference cell until it was the (6) A. Tramer, Bull. Acad. Pol. Sci., Ser. Sci., Math., Astron. Phys., 13, 751 (1965). (7) D. A. Bahniok and W. B. Person, J . Chem. Phys., 48, 1251 (1968). (8) See L. M. Julien, Ph.D. Thesis, University of Iowa, 1966, for
a review. (9) D. A. Bahniok and W. B. Person, J. Chem. Phys., 48, 5637
Experimental Section The Raman studies were carried out using a Cary The Journal of Physical Chemistry
(1968). (10) D. A. Long and D. C. Milner, Trans. Faraday SOC., 54, 1 (1958).
2311
SPECTROSCOPIC STUDYOF SOMECBr4 COMPLEXES same as in the sample cell. It was found that a practical limit to the concentration of benzene was approximately 3.0 M ; some light was still transmitted in this region by these solutions in the 0.1-mm cells. Xo reaction of CBr4 occurred during the time of the ultraviolet measurements, since the observed ultraviolet spectra of these CBr4 solutions did not change with time.
Table I1 : Molar Absorptivities "(A) at Several Wavelengths, A, and Calculated K ffor the Beneene-Carbon Tetrabromide Complex" at 25' Wavelength ,
Obsd
Calod
mfi
64A,
'oh
230 235 240 245 250 255
Experimental Results From the Ultraviolet Studies. The addition of benzene to CBr4 in cyclohexane caused an apparent increase and shift in the CBr4 absorption in the 230-265mp region, as shown in Figure 1. Six different wavelengths were chosen a t intervals throughout this region, and the absorbance measurements were made on solutions containing various benzene-CBr4 concentration ratios.
0.448 0.896 0.896 1.120 1.120 1.790 2.800 2,800
0,01823 0.01748 0.01622 0.01710 0.01628 0,01596 0.01425 0.01355
4820 4000 2650 1700 1250 1180
0,01665 0,01443 0.01447 0.01337 0.01299 0,01089 0.00868 0,00853
0.446 0.893 0.893 1.117 1.116 1.785 2.794 2.795
7770 8360 8210 6810 4920 4390
0.001737 0,002733 0,002762 0,003382 0,003526 0,004617 0.005543 0.005174
0,234 0,212 0,214 0.227 0.243 0.238 0.228 0.217 Av K f= 0.226
a The acceptor (CBr,), donor (C6He), and complex are represented by the symbols A, D, and C, respectively. The subscript T on a concentration symbol means the total concentration of that species. The concentration units are moles per liter. The molar absorptivities and concentrations are given in this table to four significant figures for purposes of consistency; the absolute accuracy with which they are determined is believed to be of the order of a few per cent, considering also systematic errors in the determination of concentration, cell length, etc.
Figure 1. The ultraviolet absorption spectra of CBrr in cyclohexane and of CBr, in benzene (cyclohexane solvent) : , 0.181 M CBr, in cyclohexane; - - - - , 0.163 M CBr, in 1.14 M benzene and cyclohexane.
The analysis of the data t o obtain formation constants and molar absorptivities was carried out using an IBR!t7040/7044 computer with an iterative method of analysis devised and programmed by Bennett." The formation constant for the one-to-one benzeneCBr4 complex is given in Table I1 together with computed extinction coefficients for the complex at the six wavelengths studied. The formation constant from these eight solutions is seen to be 0.226 l./mol. Allowing for random errors in the data, it is estimated that the probable error in this value is =t0.06, although the scatter is much less. This value for the formation constant of this complex is considerably higher than that obtained by Tramero (-0.01 l./mol). We believe that the value reported here is more accurate since our use of short path length cells allowed us to extend the study to the entire CBr4 absorption band rather than just to the beginning Of the band above 285 mp.
Identical experiments were carried out on this system at 9.5 and at 40". The formation constants computed from the data a t these two temperatures were the same (0.26 and 0.25, respectively) within the probable error as that found at 25". Evidently the heat of formation for this complex in solution is quite small, being zero within our experimental error. However, if our experimental error in Kr is actually as large as *30% (or A0.07 in Kr), then a heat of formation as large as - 3 kcal/mol could be consistent with our data, assuming a rather coincidental combination of errors. This value of Kr strongly indicates that a true complex is formed in the benzene-CBrr system. The molar absorptivities for the complexed and uncomplexed CBr4 molecules in this spectral region are shown in Figure 2. I n addition to the tc(x)values from Table 11, molar absorptivities for free and complexed CBr4 a t wavelengths higher than 255 mp are shown in Figure 2. These values for the complex were computed from the observed spectrum of CBr4 and benzene in cyclohexane using the measured Kt value. We see from Figure 2 that this band computed for complexed CBr4 is quite similar in appearance to the observed CBr4 ab(11) W. E. Bennett, unpublished results. See also the description of this method given by Julien (ref 8) and in ref 9. Volume 7.9, Number 7
July 196.9
D. A. BAHNICK, W. E. BENNETT, A N D W. B. PERSON
2312
t CBrq
6004
\
cyclohexane. I n addition, two binary solutions of CBr4 in benzene were studied by the procedure described earlier7 for the benzene-CCL system, with the 606-cm-l benzene band as the intermediate intensity standard. For all solutions, each band of interest was recorded twice, and the areas were averaged. The band areas almost always agreed within 3%. There was no evidence of a decrease in the band areas of the CBr4 bands within the time needed for the intensity measurements, indicating little or no decomposition of CBr4. From these band area measurements ( I , and Iref) the molar intensity values of CBr4were computed7n9
-TR(”) = Iv/’(IrefM)
WAVELENGTH hnp)
Figure 2. Molar absorptivities (liters per mole per centimeter) as a function of wavelength for free and complexed CBra: -_-_ , molar absorptivity for the benzene-CBrc complex; , molar absorptivities for CBr, in cyclohexane (measured and calculated).
sorption in noninteracting solution. Thus, we do not believe that we are observing a new charge-transfer band in this region, but that the local CBr4 absorption is just enhanced in the complex. However, we cannot rule out entirely the possibility that the observed intensification is due to a broad charge-transfer band which lies under the local CBr4 absorption, as suggested by Tramera6 An analogous study was also made of the tetrahydrofuran-CBrd-cyclohexane system. I n this case T H F does not absorb in the spectral region of interest; consequently, the problem of absorption in the reference cell is absent, Solutions containing various concentration ratios of T H F and CBr4 were prepared in 0.1-mm cells, and their spectra were recorded over the 320-220-mp wavelength region. Absorbance readings were made at seven equal intervals over the 230-260mp region, and these data were analyzed in the same manner as were the benzene-CBre data. The results give a formation constant value of 0.08 l./mol for the one-to-one THF-CBr4 complex. However, this value is probably not significantly different from zero12 although some intensification of the CBr4 absorption band occurred as THF was added to solution. This surprising result indicates that the oxygen-containing donor, THF, has less tendency than benzene to complex with CBr4. It is, however, consistent with the relative donor strengths of benzene and T H F toward CCL, indicated in our previous study.’ From the Raman Studies. Band area (intensity) measurements were made on the 182-, 269-, and 673cm-l bands of CBr4and on the 804-cm-l band of cyclohexane for solutions containing CBr4, benzene, and The Journal of Physical Chemistry
(1)
For this study, Irefis the band area of the 804-cm-l cyclohexane band in a solution in which cyclohexane is present to the extent of 40 vol %. I , is the integrated band area of the CBr4 Raman band displaced v wave numbers from the exciting line, and M is the molar concentration of CBr4. The [IR(”)]values for the binary benzene-CBr4 solutions were computed in a manner similar to that for donor-CCb systems.’ As I
/‘
IO1
,
00
I
I
02 VOLUME
FRACTION
I
I
0.4 OF
I
I
0.6
I
I
08
BENZENE
Figure 3. Relative molar Raman intensities of CBr, fundamentals in cyclohexane solution, also containing benzene. The curve through the points from the 269-cm-’ band was calculated assuming Kf = 0.27 and [ Z ’ R ( ~ ~ . D= ) ] ~1.66. The curve through the points from the 182-cm-l band was ~ 1.40. The calculated assuming Kf = 0.27 and [ Z ’ R ( B ~ ) ] = points for the 673-cm-1 band fall on approximately the same curve as do those shown for the 269-cm-’ band. (See text for further description of the calculation of these curves.)
were computed to in the CCl4 study,’ values of show the effect on the CBr4 band due to the added benzene, by arbitrarily taking the relative molar intensity for each CBr4 band to be 1.00 for that band in the reference solution of CBr4 in cyclohexane. These values are presented in Table 111, and the I’R(,) values plotted as a function of benzene concentration in Figure 3. (12) W. B. Person, J . Aner. Chem. Soc., 87, 167 (1966).
2313
SPECTROSCOPIC STUDYOF SOMECBr4 COMPLEXES Table I11 : Relative Molar Raman Intensities for the CBr4 Fundamentals in Solution in Cyclohexane also Containing Benzene ,---Solutionsa-[CBrrl
Bzld
IR(~~z)
IR(W
IR(6lSl
0.470 0,602 0.872 0.879 1.467
0.0 0.0 0.0 0.0 0.0
1.13f0.07 1.11 f 0.06 1.09 f 0.06 1.05 f 0.07 1.04 f 0.06
1.15 f 0.06 1.58 f 0.07 1.67f 0.07 1.67 f 0.07 1.66 f 0.07
0.84 f 0.04 0.84 f 0.05 0.85 f 0.05 0.85 f 0.04
1.451
1.114
1.21 f 0.08
1.86 =k 0.09
1.286 1.279 1.278
2.23 2.23 2.23
1.27f0.09 1.25 i 0.09 1.23 f 0.09
2.371 2.368
3.91 3.96
1.417 1.417 1.420
I’R(ISZ)
1.00 f 0.04
1.00 f 0.03
1.00 i 0.04
0.91 f 0.05
1.11 i 0.07
1.13 f 0.05
1.08 f 0.05
2.09 A 0 . 1 1.97 f 0 . 1 1.96 f 0 . 1
0.94 f 0.06 1.00 f 0.06 1.00 f 0.06
1.15 f 0.06
1.22 f 0.04
1.16 i 0.04
1.23 f 0.09 1.31 f 0.10
2.16 f 0.13 2.11 f 0.13
1.07 f 0.07 1.12 f 0.07
1.17 f 0.07
1.30 f 0.04
1.30 i 0.05
4.96 4.96 5.01
1.24 f 0.11 1.40 f 0.11 1.28 f 0.11
2.27 i 0.12 2.13 f 0.12 2.20 i 0.12
1.07 f 0.07 1.08 f 0.07 1.13 f 0.07
1.21 f 0.07
1.39 f 0.04
1.30 f 0.05
1.295 1.236 1.236
5.72 5.76 5.76
1.30 f 0.09 1.31 f 0.11 1.44 f 0.11
2.30 f 0.11 2.23 f 0.12 2.37 f 0.12
1.21 f 0.07 1.19 f 0.08 1.22 f 0.08
1.23 f 0.08
1.40 f 0.04
1.42 f 0.05
0.907 0.922
6.78 6.77
1.37 f 0.12 1.32 f 0.12
2.31 f 0.13 2.39 f 0.12
1.17 f 0.08 1.17 f 0.08
1.24 f 0.09
1.43 f 0.05
1.38 f 0.06
1.016 1.013
6.74 6.75
1.38 f 0.12 1.40 f 0.12
2.42 f 0.12 2.39 f 0.13
1.19 f 0.08 1.25 f 0.08
1.28 f 0.09
1.46 f 0.05
1.44 i 0.06
0.992 1.011
9.98 9.98
1.38 f 0.16 1.41 f 0.16
2.45 f 0.18 2.43 f 0.18
1.27f0.11 1.24 f 0.11
1.28 f 0.12
1.48 i 0.07
1.48 f 0.08
’
‘ Concentrations are in moles per liter. Defined by. eq. 1. Uncertainties are estimated standard deviations. Defined as the Benintensity of the I! band in each solution relative to its intensity in CBrr solution in cyclohexane only. The values are averaged. zene.
Discussion of Raman Results I n Figure 3 we see that the changes of I’R(”) with benzene concentration are similar for all three CBr4 fundamentals. The intensification as benzene is added is approximately the same for both the 269-cm-’ and for the 673-cm-’ bands of CBr4; the intensification of the 182-cm-’ band is somewhat less. These changes are similar to those observed for the corresponding CCl, fundamentals in solutions of CCL in benzene, except that the intensity increase did not approach a maximum value for the latter system. In the benzene-CBr4 solutions, we could not make intensity measurements on the CBr4band which corresponds to the e 218-cm-’ band of CCl,. However, the plots in Figure 3 of the Raman intensities of the CBr4 fundamentals vs. the volume fraction of benzene are obviously curved, as expected if these results are to be consistent with the conclusion from the ultraviolet study that CBr4 forms a complex with benzene. KOchanges in the frequencies of the CBr4fundamentals were observed in any of these solutions. This behavior is again similar to that of CC14 in benzene.’ These Raman data were analyzed to obtain a formation constant for a one-to-one benzene-CBr4 complex
by the method described for the analysis of the THFICN system.9 The analysis was made with the Raman intensity data for the 269-cm-’ CBr4 band which was expected to be the most accurate, since this band is strong and sharp and so the integrated intensity can be measured accurately. The values used for the final least-squares straight line plot to obtain Kf are graphed in Figure 4. From the resulting slope and intercept of the straight line shown in Figure 4,K f= 0.28 f 0.03 and [I’R(269)]c = 1.66 ;t 0.04. Here [ I ’ ~ ( 2 6 9 3 ]is ~ the relative molar intensity of the 269-cm-1 band of CBr4 in the complex. The uncertainty limits for Kf and [I’R(Z69)]0 are computed from the standard deviations in the slope and intercept of the straight line and are somewhat less than might be expected from the estimated uncertainties in Table I11 for the intensity data. However, this result from Figure 4 indicates that the Raman intensity of the 269-cm-’ band of CBr4 increases by 66% upon complexation of CBr4with benzene. The agreement of the formation constant calculated from these data with that calculated from the ultraviolet spectral results is very good and increases our confidence in this value. I n addition, the agreement between these two studies Volume 73, Number 7 Julu 1969
'
2314
D. A. BAHNICK, W. E. BENNETT, A N D W. B. PERSON strengths of the two complexes. For example, we estimated7 that the al stretching vibration was intensified by 300% for complete complexation of CC1, with benzene. This increase can be compared to the 66% intensification found here for the corresponding CBr4 vibration in the stronger complex.
'
'
'
Absolute Raman Intensities and Bond Polarizability Derivatives Bond polarizabilities and their derivatives for CBr4 in 4.0cyclohexane and CBr4 complexes with benzene can be 2.0computed in the same manner as was done previously " ' 1 4 0 00 2.0 4p 60 80 10.0M for CCL7 The standard Raman intensities [ie., @NZENEIT+ [CBrd, 45(a'J2 and 7 ( ~ ' ~have ) ~ been ] obtained by Long and Figure 4. Graph showing the least-squares determination of I\1ilnerlo for the CBr4 fundamentals from studies of K f for the benzene-CBr4 complex. Here Z* = [benxene]~/ solutions of CBr4 in CC14. We expect the standard ([z'R(Z69,]o - [Z'R(LOQ)]~), where [I'R(260)]0is the observed relative intensities of CBr, to be the same for solutions of CBr4 molar intensity of the 269-cm-' Raman line for CBr4 in a in our inert solvent, CeH12, as they are for solution in given solution and [I'R('L~O,]~ is that for the uncomplexed CBr,. the inert solvent CC1,. We have converted the standFrom this graph, the slope = 1.51 i 0.082 = ~ / ( [ Z ' R ~ Z E Q-) ] ~ [I'R(26D)]u),and the intercept = 5.33 =t0.51 = ~ / K ~ ( [ Z ' R ( L ~ Q )ard ] ~ intensities into the absolute intensity scalela used - [z'R(269)]u), where [z'R(26Q)]o is the relative molar intensity in the CC14 calculations;7 the values for both the standof the 209-cm-1 line of complexed CBr4. ard and absolute intensity scales are summarized in Table IV, together with the values of the polarizability derivatives, a f t and y f z , for each mode. I n addition, indicates that the same chemical species causes the we have listed in Table IV the absolute intensities of observed changes in both the ultraviolet and Raman the CBr4 fundamentals for the benzene-CBr4 complex spectra. which we calculated from the absolute intensity values Using this value of the formation constant (0.28) the for uncomplexed CBr4 by multiplying them by the relative molar intensities for the other bands of comintensification factors reported above for the complex. plexed CBr4 can be calculated. Thus it is found that Since the intensity change of the 128-cm-l CBr4 band the 182-cm-l band is intensified by 40% (=k7%) in the could not be studied, we have assumed that it has the complex and the 673-cm-' band intensified by 65% same behavior on complexing as did the analogous (AS%). The curves through the points shown in Figure 218-cm-' band of CCl, which did not change as 3 have been computed using these values of the intensibenzene was added to solutions of CCL ties and of the formation constant; the agreement beBond polarizabilities and their derivatives are caltween the computed curves and the observed points is culated by the same procedure as for CCL7 The bond quite satisfactory. polarizability parameters in the zero-order approximaThe Raman intensity changes observed for the CCll tion14 are related to the experimental values of the fundamentals in various aromatic and amine solvents7 derivatives with respect of the isotropic (a'%)and anisoare similar to those observed here for CBr4 in benzene. tropic (y'J parts of the polarizability. From the A detailed comparison can be made for the al and f bond polarizability parameters, atnl and cy',? (the symmetry vibrations of the two tetrahalide molecules. derivatives of the parallel and perpendicular comFor each molecule, the a1 and f symmetry stretching ponents of the bond polarizabilities aril and annwith vibrations show an intensity increase of the same respect to the bond distance for the nth bond) can be magnitude for solutions of the tetrahalide in benzene. c ~mputed.~ Furthermore, the f bending mode increases but by a The values of bond polarizabilities and their derivasmaller amount. The similarity in intensity behavior tives have been calculated, and the results are presented suggests that the same type of interaction is responsible in Table V. For the two-by-two f symmetry block in for the observed changes of the fundamentals of both the computation of y f C - B r we have used L matrix comCX4 molecules measured in benzene solutions. Howputed from the calculation by Pistorius and Haarhoff .15 ever, the nearly linear changes (as the concentration of The calculated values of the bond polarizability donor changes) in the Raman intensities observed for parameters which are listed in Table V for CBr4 in the cc14 fundamentals suggest that the formation constants for the donor-CCl4 complexes are quite small (13) G. W. Chantry and R. A. Plane, J . Chem. P h y s . , 32, 319 (1960). (less than about 0.1 l./mol). However, when we com(14) iM.Wolkenstein and M. Eliashevioh, Acta Physicochim. U R S S , pare quantitatively the magnitude of the intensifica2 0 , 525 (1945). tion of the Raman bands for the two tetrahalides, we (15) C. Pistorius and P. Haarhoff, 2. Phys. Chem. (Frankfurt ani see that this intensification is not proportional to the Main), 19, 202 (1959).
&,-
The JouTnal of Physical Chemistry
SPECTROSCOPIC STUDY OF SOME CBr4 COMPLEXES
2315 ~
~~~~~~
Table IV : Standard and Absolute Intensities of the CBr4 Fundamentals of CBr4 in CC14 and of CBrh Complexed with Benzene 7(r’d2
7(~‘a)~
1.15
0.069
0.720
23.29
1.397
14.58
38.66
1.397
24.06
45(a’1)~
Standard intensity” (CBr4 in CC&) Absolute intensityb (CBr4 in CCh) Absolute intensity’ (CBr4 complexed with benzene)
7 (7‘4)s ,
P’1
Y’2
Y‘8
Y ’4
2.916
0.7194
0.1996
1.443
0.6454
4.082
0.9269
0.1996
1.854
0.7635
0.144
Values are from Long and Milner.10 The subscripts on a and y refer to the numbering of the vibration given in Table I. The Calculated by multiplying the standard intensity standard intensities are estimated by Long and Milner to be accurate to *lo%. Calculated by multiplying the absolute intensities for the CBrl mode for CBr4 in values by 20.25 A4 (awu)-l (see ref 7 and 12). CC14 by the values of Z’RCgiven in the text. The value of 7(y’# (from the 128-cm-l band) is assumed not to change from uncomplexed CBr4 (see text).
’
cyclohexane show that the zero-order bond polarizability approximation fits the intensities of uncomplexed CBrd quite well, since the value of Y C - B ~ computed from ylZagrees with that computed from 7’3 and $4. This result is probably fortuitous because of the uncertainties in both the intensities and especially in the L matrix. The zero-order approximation also holds quite well (within our experimental uncertainty) for the complexed CBr4. We note with some interest that complexing the CBr4 molecules causes an increase in all the bond polarizability derivatives, except possibly for YCBr. This change upon complexing is consistent with the changes we have observed in our previous studies7~g and such increases appear to be characteristic of donoracceptor interactions, at least. It remains to be seen whether such increases in intensity are always to be expected for all kinds of complexes. Perhaps the most striking conclusion from this study is the verification of the existence of a very weak benzene-CBra complex in solution. The existence of this
Table V : Bond Polarizability Parameters for CBr4 in Cyclohexane and for Benzene-CBr4 Complexes
CBr4 in cyclohexane (CC14) CBr4 complexed with benzene
a’c-Bn
Y‘C-~r,
A2
A1
9.64a 12.42
YC-Br,
AB
a‘m, A2
A2
6.38
2.65’
7.47
1.09
7.74
2.65‘ 2 . 9Pjd
9.30
1.56
This value is to be compared with the value ( . I ’ ~ - ~=~9.99) reported by T. V. Long, 11, and R. A. Plane, J. Chern. Phys., 43, 457 (1965), measured in an unknown solvent. Note that the conversion from the value listed in their Table I1 to our intensity scale is made by multiplying their value by 3.0. The same value is calculated from y’l and from y l S , y ’ ~(see ref 7). ‘ Calculated from 7’2 for complexed CBr4, assuming this parameter (from the intensity of the 128-cm-l band) does not change from that for the uncomplexed species. Calculated from y‘a and y’r of the complexed CBr,.
’
complex tends to strengthen the belief that the interaction observed between benzene and CCI, (see ref 7 ) is really due to incipient complex formation, even though the formation constant of that particular “complex” is too small to measure.
Stability of CBr4Complexes The formation constant for benzene-CBra (0.27 l./mol) is comparable to that of benzene-12; however, the stability, as measured by - - H I , the enthalpy of formation, is apparently somewhat less than for benzene-12. If we consider the entropy of formation for these two complexes, we see that such an effect may be expected. Thus, the formation of both complexes involves a loss of symmetry. For 1 2 the symmetry number of the free molecule is two; for benzene, it is twelve; and for CBr4, it is twelve; in an axial Bz-12 complex the symmetry number is only six; while it is three for the CBr4 complex. Hence, if the interaction between benzene and each acceptor is otherwise similar, the entropy of formation of the CBr4 complex may be expected to be more positive than that for the I2 complex by R(ln 12) or about 4.94 eu. Thus, for approximately the same formation constants (or free energy of formation) at room temperature, the enthalpy of formation for the CBr4 complex may be expected to be less negative than that of the IZcomplex by about 1.5 kcal/ mol. Hence, -AHf for the benzene-CBr4 complex is expected from this argument to be approximately zero. The stabilization of a complex which is as weak as benzene-CBr4 is very possibly caused by electrostatic forces alone. l6 The observed increased stability of the benzene-CBra complex over that of benzene-CClr can then probably be explained as due to the higher polarizability of CBr4 as compared to CC1,. It does not seem reasonable to assume that the electron affinity of CBr4 is so much larger than that of CC14 that it can account for the increased stability of the former complex (16) For example, see M. W. Hanna, J. Amer. Chem. SOC., 90, 286 (1968).
Volume 73, Number 7
J u l y 1969
H. E. SPENCER
2316 as a result of increased charge-transfer resonance stabilization. We might further expect that the electrostatic interaction might well vary in magnitude as the C-Br bond stretches in a vibration and so cause the observed Raman intensification. This hypothesis that the interaction between benzene and CBr4 may be mainly electrostatic is also consistent with the observed change in the ultraviolet spectrum, since the latter appears to be just a shift of a locally excited CBr4 transition. Furthermore, it explains why the interaction between the stronger electron donor (THF) and CBr4 is found to be weaker than that between benzene and CBrd, and it is consistent with the small amount of charge transfer found in the solid xylene-CBr4 complex by the quadrupole moment s t ~ d i e s . The ~ fact that the vibrational frequencies of the CBr4 molecules do not
appear to change in the complex suggests either that there is very little charge-transfer action, or else that the acceptor action on CBr4 is not into the n* C-Br orbital (in contrast with the acceptor action of Brz, for example). In summary, we favor, at present, the hypothesis that the benzene-CBrr interaction is primarily electrostatic, although we do not eliminate the possibility that some very weak charge-transfer action may occur. Aclmowledgments. We are grateful to the Public Health Service for financial support through Grants GM-10168 and CAI-14648. The authors are also indebted to the University of Iowa Computer Center and Graduate College for assistance with the computing costs.
Quantum Yields of Production of Iron(I1) and of Cobalt(I1) in the
Photolysis of Solid Potassium Trisoxalatoferrate(111) Trihydrate and Potassium Trisoxalatocobaltate(111) Trihydrate’ by H. E. Spencer Research Laboratories, Eastman Xodak Company, Rochester, New York 14660 (Received November 22, 1968)
The values of quantum yields of production of Fe(II), GFe(II), are about an order of magnitude lower in the microcrystalline solid compared with solution values, ranging between about 0.07 and 0.17 and increasing with increasing intensity and with decreasing wavelength. Values of GCo(I1) in the microcrystalline solid vary between about 0.07 and 0.55, are somewhat less than values for solutions, and are independent of intensity, but decrease with increasing exposure at the shorter wavelengths. It is suggested that hole-electron pairs are generated within the crystals; these subsequently generate radical ions at the surface.
Introduction The photochemistry of solutions of potassium trisoxalatoferrate (111) trihydra te, Ka [Fe(C,O,) 3 ] * 3Hz0, and potassium trisoxalatocobaltate(II1) trihydrate, K3[ C O ( C ~ O.~3Hz0, ) ~ ] has been studied extensively.2-6 Far less work has been devoted to the photochemistry of the solids themselves. In fact, the only quantitative studies are those of Wendlandt and Simmons in which they determined the stoichiometries of the photolysis reactions.6,’ I n both solids, as well as in solutions of them, the tripositive metal ion is reduced to the dipositive metal ion, and oxalate is oxidized to carbon dioxide. Also in the solid, a color change occurs upon photolysis : the light green iron(III) compound turns dark yellow, and the dark green cobalt(II1) compound turns pink. The Journal of Physical Chemistry
It is the purpose here to report the solid-state quantum efficiencies of production of Fe(I1) and Co(II), (1) PFesented at the 156th National American Chemical Society Meeting, Atlantic City, N . J., Sept 1968. (2) (a) C. A. Parker, Proc. Roy. Soc., Ser. A, 220, 104 (1953); (b) C. G. Hatchard and C. A. Parker, ibid., 235, 518 (1956); (0) J. G. Calvert and J. N. Pitts, Jr., “Photochemistry,” John Wiley & Sons, Inc., New York, N. Y., 1966, pp 783-786. (3) G. E. Porter, J. G. UT.Doering, and S. Karanka, J . A M . Chem. SOC.,84, 4027 (1962). (4) A. W. Adamson and A. H. Sporer, ibid., EO, 3865 (1958); J. Inorg. Nuel. Chem., E, 209 (1958). (5) T. B. Copestake and N. Uri, Proc. Roy. SOC.,Ser. A , 228, 252 (1955). (6) W. W. Wendlandt and E. L. Simmons, J . Inorg. Nucl. Chem., 28, 2420 (1966). (7) W. W. Wendlandt and E. L. Simmons, ibid., 27, 2317 (1965).
