D. FROST,F. HERRING, A. KATRIB, C. MCDOWELL, AND R. MCLEAN
1030
signment of the structure on the electronic absorption band as the result of vibronic transitions involving the anthracene ring symmetric stretching mode. Of course, the -1300-crn-’ value would be characteristic of the mode for the electronic excited state. D . Conductivity Results. Ubbelohde, et al., have described rather extensive studies of the electrical characteristics of the alkali metal-anthracene crystals prepared from solution. They conclude that the conductivity increases significantly as the metal to anthracene ratio is increased, but that resistivities remain quite high-in the lo8to 10’0 ohm-cm range at 20”. Though our thin film spectroscopic results seem to contradict their view of the solids as solutions of metal atoms in anthracene with a continuously variable composition, the new data support their contention that electron
transfer occurs from the metal atoms to anthracene. Further, our resistivity data, obtained by the four-point probe method, are reasonably compatible with the published values from the bulk samples. Thus, the resistivity was observed to decrease roughly two decades as the metal to anthracene ratio was varied from 1 : 2 to 2: 1. However, the measured thin film resistivities were significantly lower than those reported for bulk samples of comparable composition, the range extending from lo6 to los ohm-cm a t -100°K. Our data do not warrant a more quantitative presentation at this time.
Acknowledgments. This research was supported by the National Science Foundation under Grant GP24256, thc Paint Research Institute, and the Oklahoma State University Research Foundation.
Photoelectron Spectra of CH,SH, (CH,) 2S, C,H,SH, and C,H,CH,SH;
the Bonding between Sulfur and Carbon by D. C. Frost, F. G. Herring, A. Katrib, C. A. McDowell,* and R. A. N. McLean Department of Chemistry, University of British Columbia, Vancouver 8 , British Columbia, Canada (Received November 17, 1971) Publication costa assisted by The National Research Council of Canada
The first ionization potential in the photoelectron spectra of CH3SH (9.42 eV) and (CHJzS (8.65 eV) as in H2S (10.47 eV) corresponds to ionization of a lone-pair electron, and, although there is a small T interaction in the C-S bond, the inductive effect is thought to cause the considerable shift to lower energy. I n contrast to these molecules and a-toluenethiol (first I.P., 9.25 eV) the sulfur lone pair in thiophenol exhibits a considerableamount of T interaction with the adjacent benzene ring. The higher I.P.’s of CH3SH and (CHJZShave been assigned, and it is shown that ionization occurs more readily from u orbitals situated mainly in the C-S bond than the S-H bond.
Introduction Photoelectron spectroscopic studies have been carried out previously on divalent sulfur compounds.1p2 I n particular the PE spectrum of hydrogen sulfide is the subject of intensive research.2 The present study reports the effect on thc PE spectrum of successive replacement of the hydrogen atoms in H2S by alkyl and phecyl groups. Divalent sulfur is generally thought to participate in bonding via tu70 u bonds involving mainly two of the three 3p orbitals on the sulfur atom leaving the 3s orbital and the remaining 3p orbital to accommodate the four nonbonding electrons. The possibility of the sulfur 3d electrons playing a role in the bonding is disThe Journal of Physical Chemistry, Vol. 76. N o . 7 , 1972
cussed in section D later in this paper. The formally “3s nonbonding pair” is much more tightly held by the sulfur atom than the “3p nonbonding pair” and is directly involved in molecular bonding. The replacemcnt of hydrogen atoms in HzS by electron-donating groups leads to a destabilization of these electron pairs, which is reflected in the ionization potentials of these molecules. The first ionization potentials of a number (1) W. R. Cullen, D. C. Frost, and D. A . Vroom, Inorg. Chem., 8, 1803 (1969). (2) (a) J. Delwiche and P: Natalis, Chem. Phys. Lett., 5, 564 (1970); (b) D. W. Turner, C. Baker, A. D. Baker, and C. R. Brundle, “Molecular Photoelectron Spectroscopy,” Wiley, New York, N. Y . , 1970, Chapter 4; (c) D. C. Frost, A. Katrib, C. A. McDowell, and It. A . N. McLean, Int. J.Mass Spectrom. Ion Phys., 7,485 (1971).
THEBONDING BETWEEN SULFURAND CARBON
1031
Table I : Vertical Ionization Potentials of Methylmercaptan and Dimethyl Sulfide (eV) 7
CNDO/Z
8.4 10.8 14.5 17.1 17.3 21.4
a
Methylmercaptan Experimental
9.42 12.0 13.9 15.0 15.5 -20.0
From ref 1 (adiabatic I.P.’s).
Orbital type
CNDO/2
n(S3p) a ” u(S-C) a’ u(S-H) a‘ u(CH3) a” u(CHa) a” n(S,3s) a’
7.7 11.4 11.5 15.1 16.1 17.5 18.4 20.3
---- Dimethyl sulfide-----Experimental
Orbital* type
8.68” 8.65 10.96 11.2 12.16 12.6 13.68 14.2 14.8 15.4 15.7 19.7
’Assuming CZ, symmetry.
of divalent sulfur compounds earlier obtained by other methods3 illustrate this effect, e.g., CH3SH (9.44 eV), C2H5SH(9.285 eV), C2H5CH2SH(9.195 eV), CH3SCH3 (8.685 eV), C2H5SCH3(8.55 eV), and C2H6SC2H5 (8.43 eV). Obviously the larger the alkyl group, the greater will be the destabilization and the lower the ionization potentials. We wish to report here the high-resolution 584-A photoelectron spectra of methyl mercaptan, dimethyl sulfide, thiophenol, and a-toluenethiol together with CNDO/2 calculations which are used to aid the interpretation in the absence of more rigorous calculations. Experimental Section The photoelectron spectrometer used in this work will be fully described else~vhere.~The energy source is the helium I 584-A (21.22 eV) resonance line. Calibration of the spectra was accomplished by inclusion of small amounts of argon as an internal standard. The sample pressure in the instrument was about Torr. The materials for the study were obtained from Matheson Co. Led. (CH3SH) and K & K Laboratories, Inc. ( (CH3)2S,C6H5SHand CeH5CH2SH). There were no obvious impurity lines observed in the spectra. In the following discussion Koopman’s theorem is assumed to hold. The vertical ionization potentials and results of CND0/2 calculations for CH3SH and (CH&S are listed in Table I. Results and Discussion A . “Nonbonding” Electrons. It is to be expected, as in HZS, that the first electron removed from molecules of the type RSH or R2S will come from the “3p lone-pair” orbital unless one of the groups R possesses energy levels of its own near the energy of the sulfur 3p orbital. This situation occurs in thiophenol and a-toluenethiol where the benzene ring I.P.’s lie at about the same energy as the sulfur lone-pair I.P. The PE band produced on ionization of a nonbonding electron is characteristically sharp reflecting the negligible change in the Franck-Condon envelope between the neutral species and the ion. Consequently the lone-
L (CH& S
85
9
950
100 eV
Figure 1. First I.P.’s of HZS, CHsSH, CeH.&H2SH, and (CH3)nS observed in their photoelectron spectra obtained using the He 584-A resonance radiation.
pair I.P.’s in CH3SH and (CH3)*Sare readily identifiable, as can be seen from Figure 1, at 9.42 and 8.65 eV, respectively, compared to 10.47 eV for H2S. These compare well with earlier values.386 Therefore, as expected, the sulfur lone pair has become progressively destabilized. The CNDO/2 calculations6 we have carried out predict the separation of the lone-pair ionization of these compounds (HzS, 9.4; CH3SH, 8.4; (CH&S, 7.7 eV) well, although after the usual 4-eV scale has been employed the predicted I.P.’s are about 1 eV too low. The results of the calculation indicate that the methyl group orbitals destabilize the sulfur 3p (3) K. Watanabe, T. Nakayama, and J. Mottl, J. Quant. Spectrosc. Radiat. Transfer, 2 , 369 (1962). (4) D. C. Frost, A . Katrib, and C.A. McDowell, t o be published. ( 5 ) G. Hersberg, “Electronic Spectra of Polyatornic Molecules,” Van Nostrand, Princeton, N . J., 1967. (6) J. A. Pople and D. L. Beveridge, “Approximate Molecular Orbital Theory,” McGraw-Hill, New York, N. Y., 1970.
The Journal of Physkal Chemistry, Vol. 76, N o . 7, 1979
1032
D. FROST, F. HERRING, A. KATRIB,C. MCDOWELL, AND R. MCLEAN
orbital and produce an increasingly S-C a-antibonding molecular orbital. The vibrational spacing evident in the photoelectron spectra should give some indication of this antibonding charactcr. For CHSSH, the vibrational spacings are 680 40 and 1250 80 em-’. The first obviously corrcsponds to the C-S stretching vibration, which is expected to be cxcitcd if there is any interaction between thc sulfur p a orbital and a carbon p orbital. As the C-S strctching frequency is 704 crn-I5 in the neutral spccics, it is impossible to say whcthcr thc highest orbital is bonding or antibonding with rrspect to the carbon-sulfur bond, but whichever it is the interaction is apparently fairly small. Tlic 1250-cm-I vibration almost crrtainly cor1 rsponds to the symmetrical dcformation of the CHs group (reduced from 1335 cm-1 in thc ricutral spccics) in line with the slight amount of CH, u-bonding charactrr in the. mainly lone-pair orbital. It would be uscful to compare the vibrational structurc on the first PES band of CH3SH with that of thc alkyl halides. The best molcculc for such a comparison would bc CHJI n-hcrc thc spin-orbit coupling is large enough to produce two components 1 ~ 1 1enough scparatcd such that their mutual interaction is small. Also as far as intcraction of the “lonc pair” with othcr orbitals is concwncd, it is best to comparc sulfur and iodinc, since CH31 has its first I.P. a t ~ 9 . 5 cV.7 0 In CHJ thc principal vibration excited is v2, the symmctric dcforrnation (1290,21260,8or 1225 cm-I ’). The C-I strctching vibration i5 also cvidcnt,7~8though it is not wcll rcsolvcd in the former study, and the symmetric CH-strrtching vibration has bcen observed7 as well. In CH3I there is similar slight reduction in the C-I strrtching frclqucncy on ionizing, indicating a slightly a-bonding interaction in thc C-I bond. It seems thcrcforc that if anything the slight rcduction for CH3SH may just bc significant despite the error limit and that in contrast to thc calculation there is a very small amount of a-bonding character in the C-I bond. Howcver as a whole the sulfur lone pair is destabilized comparcd to that of H2S by interaction with the methyl group, mainly by u donation. The substitution of another H atom by a CH3 group does not havc as large an effect on the I.P., but the vibrational frequency is now somewhat greater for the ion (750 f 60 cm-l) than in the neutral species (685 ~ m - l ) . ~The lone pair must bc w r y slightly antibonding with respect to the two carbon atoms which must be bonding with respect to cach other. This was also shown in thc Cr\;D0/2 calculations. It seems that the large shift of thc first I.P. to lower energy in (CH3)2S arid CH3SH compared to H2S is due principally to a large u-inductive effect raising the coulomb energy of the sulfur 3p orbitals, and to a much lesser extent, to a intcraction between sulfur and carbon orbitals. The a interaction of a sulfur lone pair with methyl groups is relatively small. However, we shall now see
*
*
T h e Journal of Physical Chemistry, Vol. 76, 1\70.
7 , 1972
that the opposite is the case when sulfur is attached to a phenyl group. This can best be seen by comparing the PE spectra of benzenethiol (PhSH) and toluenethiol (PhCH2SH). In benzene the lowest I.P. due to ionization from the a elg orbital is a t 9.25 eVIO and in monosubstituted benzenes (C2, or lower symmetry) this is split into two orbitals, one having a R interaction with the substituent (bl in C2,) and one non- (or less) intciacting (az in (22,). The separation between these two levels is greatest with substituents with polarizable lone pairs and least for a substituent with no lone pairs.” Toluene is of this latter type and the PE band exhibits only a slight broadening and a shift to lower energy compared to benzcne.11*12Replacement of one of the H atoms of the CHHgroup by the SH group produccs no remarkable changes in the PE spectrum except for the superposition of a new sharp peak (at 9.25 eV) mith a t least one other vibrational component (spacing -1220 cm-l). The first vertical I.P. is 8.85 eV (cf. 8.81 eV in toluene). It can be seen in this case that the sulfur lone pair has an almost negligible effect on the orbitals of the benzene ring. The inverse is also true, as besides the sharp band, PhCHzSH has its sulfur lonepair I.P. a t 9.25 eTT (Figure 1) compared with that of CH3CH&H a t 9.285 eV.3 The vibrational spacing of 1220 f 50 em-’ probably corresponds to the wagging motion of the CH2 group a t -1400 cm-I, as therc will be some interaction between the sulfur lone pair and the CHZ orbitals. I n contrast to a-toluenethiol, the replacement of one H atom in benzene by a SH group has an enormous effect. Below 11 eV instead of the one I.P., therc are three a t vertical I.P.’s 8.28, 9.38, and 10.65 eV, respectively (Figure 2 ) . The 9.38-eV band obviously corresponds to the noninteracting benzene ring orbital. The separation between the other two I.P.’s (2.37 eV) shows the size of the interaction between the sulfur lone pair (originally a t -9.4 eV) and the ?r (“B1”) orbital of the berizcne ring (originally a t 9.38 eV). Iodobenzene also shows a large interaction, but the separation between the two bl I.P.’s is only 1.78 eV and the lowest I.P. is a t 5.67 eV.” This latter is also mainly a (C&) as the az a (C6H5)is still a t 9.38 eV and the bz a (I)is a t 9.64 eV. The 8.28-eV I.P. of PhSH is slightly narrower than the 10.65-eV band and because the 0 + (7) J. L. Ragle, I. A. Stenhouse, D. C. Frost, and C. A. McDowell, J . Chem. Phys., 53, 178 (1970). ( 8 ) A. W. Potts, H. J. Lempka, D. G. Streets, and W. C. Price, Phil. Trans. Roy. SOC.London, Ser. A , 268,59 (1970). (9) G. Herzberg, “Electronic Spectra of Polystomic Molecules,” Van Nostrand, Princeton, N. Y., 1967, p 652. (10) 1. D. Clark and D. C. Frost, J . A m e r . C‘hem. Soc., 89, 244 (1967). (11) (a) D. W’.Turner, C. Baker, A . D. Baker, and C. R. Brundle, “Molecular Photoelectron Spectroscopy,” Wiley, New York, N. Y., 1970, Chapter 11; (b) A. D. Baker, D. P. May, and D. W. Turner, J . Chem. SOC.B, 22 (1968). (12) A. Cornford, D. C . Frost, and R. A. N. McLean, unpublished results.
THEBONDING BETWEEN SULFUR AND CARBON
1033 CH,SH
b
8
9
10
11 e V L
Figure 2. The photoelectron spectrum of C8H6SHobtained using the He 5 8 4 - i resonance radiation. +
IO
0 transition is the vertical one, we believe the corresponding orbital to have more sulfur lone-pair character. The vibrational structure on the band consists of a progression of eight peaks separated by -400 cm-l, which probably corresponds to the substituent-sensitive ring breathing mode at 412 cm-’ in the neutral thiophenol. l 3 CND0/2 calculations carried out on a-toluenethiol assuming the C(H2)-S-H plane to be perpendicular to the ring predict the second ionization potential to be associated with the removal of an electron from the sulfur lone-pair orbital. This lone-pair orbital is almost completely noninteracting as regards the benzene ring. Slight deviations in this configuration are predicted to cause little change in the ionization potential. However, if the C(H,)-S-H plane coincides with the plane of the benzene ring and becomes the origin of the lowest I.P. comparison of the calculations and photoelectron spectra indicates that the former configuration is preferable. In the case of thiophenol the calculations show that the first I.P. arises from an orbital which is formed mainly from the sulfur lone pair with a considerable amount of benzene ring character, which is in agreement with the interpretation of the PE spectrum suggested above. B. Sulfur-Carbon and Sulfur- Hydrogen a-Bonding Electrons. For methanethiol and dimethyl sulfide, the second and third ionization potentials (see Figures 3 and 4) should correspond to ionization from mainly the C-S-H and C-S-C bonding regions, respectively. I n CH3SH the second and third I.P.’s are at 12.01 and 13.9 eV, neither band having any resolvable vibrational structure (the lack of structure in the second band is probably due to a dissociation to CHBS+a t 11.2 i 0.5 eV).14 By comparison with HzS and (CHs)2Sit should be possible to decide which orbital has more S-H character and which more s-C character. I n H2S the S-H bonding orbitals are at -13.4 (al) and -15.2 eV (bz), and in (CH&S we assign the C-S bonding orbitals I.P.’s as 11.2 (al) and 12.6 eV (bz). Thus, by taking the average of these two it seems that the 12.0-eV band arises from an orbital with more S-C character and the 13.9-eV band from an orbital with more S-H character. The CNDO/2 calculations show an almost equal amount of each with the second I.P. MO having most
12
14
16
eV
Figure 3. The photoelectron spectrum of CHZSH obtained using the He 584-1 resonance radiation.
of its density on the sulfur and the third on carbon and the unique hydrogen, hence the usual16 overestimation of the separation between them. Methyl iodide (which we have compared with methanethiol above as regards the lone pairs) has its C-I a-bonding I.P. at 12.50 eV, which is again fairly close to the mainly C-S a-bonding I.P. at 12.0 eV. C. Other Bands of Methanethiol and Dimethyl Sulfide. As in the alkyl h a l i d e ~ ~ *the * , ’ ~higher energy I.P.’s below 21.22 eV can be assigned as from mainly CH3-bonding orbitals. Around 15 eV in the alkyl halides there is a band with a Jahn-Teller contour assigned as ionization from the e (CH,) bonding orbital.*Vle I n methanethiol there is a broad band of intensity twice that of the 12.0- and 13.9-eV bands with twin maxima a t 15.0 and 15.5 eV which is almost the same as occurs in the PE spectrum of CH3Br. It could almost be considered that the CH3 group is only observing the SH group as a single atom, and certainly the CYD0/2 calculations indicate this as the two low-energy bonding CHa I.P.’s are almost degenerate at 17.11 and 17.27 eV. The other CHs-bonding orbital which appears to be almost all C(2s) produces a very diffuse band in the alkyl halides PE spectra at -20 eV, and the same thing is observed in the case of methanethiol. For dimethyl sulfide four Iow-energy CH3-bonding orbitals are expected and in the broad band from 13.50 to 16.4 eV four components can be isolated. The CNDO/2 calculations indicate a much larger than observed range for these CH3 orbitals and probably over(13) D. W. Scott, J. P. McCullough, W. N . Hubbard, J. F. Messerly, I. A . Hossenlopp, F. R . Frow, and G. Waddington, J . Amer. Chem. Soc., 78,5463 (1956). (14) W.E. W. Ruska and J. L. Franklin, Int. J . Mass Spectrom. Ion Phys., 3, 221 (1969). (15) We have generally found that CNDO/2 calculations tend to
overefitimate the interaction between symmetry related orbitals of similar energies. This can lead t o errors in the assignment of PE spectra, but with consideration of this complication this can be eliminated at least in small molecules. (16) D. W. Turner, C. Baker, A. D. Baker, and C. R. Brundle, “Molecular Photoelectron Spectroscopy,” Wiley, New York, N. Y . , Chapter 8.
The Journal of Phgsical Chemistry, Vol. 76, N o . 7, 1973
D. FROST, F. HBRRING, A. KATRIB, C. MCDOWELL, AND R. MCLEAN
1034
c
1 0
12
14
16
18
20
Figure 4. The photoelectron spectrum of (CH,),S obtained using the He 584-A resonance radiation.
estimate the interaction between orbitals on the two CH3 groups. D. Sulfur Sd-Orbital Participation. There has been considerable discussion on the involvement of 3d orbitals in the bonding of bivalent sulfur compounds.17-1s It is possible to make deductions about d-orbital participation in the normal valence states of second row elements from the PE spectral data,20although there may be other interactions which produce the same effects. We believe some of our results are pertinent in this regard. Goodman and TaftZ1explained the decrease in the intensity of the ‘Lb ‘A transition in going from thiophenol to p-methylthiophenol in their substituent interference experiments in terms of a strong interaction between the A orbitals of the benzene ring and sulfur 3d orbitals. However, remeasurement of the spectra by Di Lomardo and ZauliZ2indicated that there was a slight increase in intensity from thiophenol to p-methylthiophenol. They inferred that there is little involvement of the sulfur 3d orbitals and that the only perturbation of the benzene ring A orbitals .is via S(3p?r)-C(2pa) bonding. This latter interaction has been considered fully above, but we also have evidence to support Di Lomardo and Zauli’s conclusion. Again assuming Czvsymmetry for thiophenol, it is possible for the “az” benzene ring A orbital to interact with the sulfur d,, orbital and the “bl” benzene A orbital to interact with +
The Journal of Physical Chemistry, Vol. 76, N o . 7, 197.3
the sulfur d,. orbital. In the first case the overlap will be very small as the constituent orbitals involve nonadjacent atoms, and, as expected the [‘az’’?r orbital occurs a t almost exactly the same energy ( h 0 . 1 eV) as in PhOH, PhNHz, PhCH3, PhSiH3, and many other monosubstituted benzenes. The real test of sulfur d-orbital interaction is in the “bl” ?r orbitals with I.P.’s a t 8.28 and 10.65 eV, respectively, which, as mentioned above, must have almost equal amounts of benzene A and sulfur lone-pair character. If the unperturbed benzene ring bl orbital I.P. were at the average of these two I.P.’s i.e., 9.46 eV, the stabilization compared to the az orbital would be negligible. It could be explained by the expected greater inductive effect of the sulfur atom on the bl orbital. As both levels would be almost a t the same energy if they were unperturbed -9.40 eV, it is difficult to find any other method of measuring the interaction. We believe that this indicates if there is any stabilization of the bl ?r orbitals by the use of sulfur 3d orbitals, it is very small and has little effect on the orbital energies. The main effect on the orbital energies is the interaction between the sulfur 3pn orbital and the benzene ?r orbital. Acknowledgments. We are grateful for generous financial support from the Kational Research Council of Canada. One of us (R. A. K. 11.)thanks the University of British Columbia for a Killam Fellowship. (17) D.R.Williams and L. T . Komtik, J. Chem. Soc. B , 312 (1971), and references therein. (18) M.J. S.Dewar and N. Trinajstic, J.Amer. Chem. Soc., 92, 1453 (1970). (19) D. P. Santry and G. A. Segal, J. Chem. Phys., 47, 158 (1967); D. P. Craig and T . Thirunamachandran, iM., 45, 4183 (1965); E. A. C. Lucken, J. Chem. SOC.A , 991 (1966); E. A. C. Lucken, Thew. Chim. Acta, 1 , 397 (1963); I. 13. Hillier and V. R. Saunders, Chem. Commun., 1183 (1970). (20) D. C. Frost, F. G. Herring, A. Katrib, K.A. N. McLean, J. E. Drake, and N. P. C. Westwood, Chem. Phys. Lett., 10, 347 (1971); Can. J . Chem., 49,4033 (1971). (21) L. Goodman and R. W. Taft, J. Amer. Chem. Soc., 87, 4385 (1965). (22) G.Di Lomardo and C. Zauli, J. Chem. SOC.A , 1305 (1969).
