J . Phys. Chem. 1987, 91, 3969-3974
3969
Without detailed calculations or further experimental work it is not possible to say with certainty that the N(CH3)3line widths are dominated by homogeneous broadening. However, it seems unlikely that the line widths obtained for N(CH3), are entirely due to inhomogeneous broadening. The line widths are far larger than those estimated from ground-state rotational constants and are larger than line widths for As(CH,)~and P(CH3)3which do correlate well with estimates based on ground-state rotational constants. Thus it is more likely that observed line widths obtained for P(CH3)3and As(CH3), are largely inhomogeneous in origin with perhaps some homogeneous contribution (especially where resonances are expected). The predominance of inhomogeneous broadened line widths for P(CH3), and As(CH3), could be confirmed (or denied) by measurements of line widths as a function of temperature to determine if the changes in rotational contours are consistent with changes in rotational populations.
lecular symmetry plane containing the C, M, and H b atoms ( M = N, P, or As); (b) transitions which correspond to C-Hb bond absorptions; these transitions involve C-H bonds in the molecular symmetry plane and are at lower energy than the C-H, absorptions; and (c) combination bands corresponding to local-modenormal-mode transitions. Overtone studies of N(CH,),, P(CH,),, and As(CH3), indicate that the trans effect, involving lone pairs of electrons, previously found in compounds with elements in the first and second row, also operates in the molecule As(CH,),, which contains a third-row heteroatom. It appears that the overtone line widths of P(CH3), and As(CH,), can be explained principally via inhomogeneous mechanisms. For N(CH3)3the overtone line widths have a significant component due to homogeneous broadening. This change in the broadening mechanism appears to correlate with the M-C bond length in the M(CH3), compounds.
Conclusions
Acknowledgment. We gratefully thank Dr. R. N. Jones and his collaborators from the National Research Council of Canada for the computer programs used in the deconvolution of the spectra. This work was supported by the National Science Foundation under grants CHE82-06976 and CHE85-06957. Registry No. N(CH3),, 75-50-3; P(CH,),, 594-09-2; As(CH,),,
The spectra of C-H overtones of N(CH3),, P(CH3),, and A s ( C H ~ have ) ~ been investigated by using intracavity photoacoustic spectroscopy and standard infrared techniques. Three. types of vibrational transitions are identified for C-H absorptions of these molecules: (a) transitions which correspond to C-Ha bond absorptions; these transitions involve C-H bonds out of the mo-
593-88-4.
Electron-Donating Ablllty of Ethyl and Ethenyl Groups from Core-Electron Spectroscopy and ab Initio Theory. A Study of CH3CH,X and CH,CHX (X = F, CI, Br, I ) M. R. F. Siggel:* G. S. Nolan: L. J. Saethre,**T. D. Thomas,*+ and L. Ungiert Department of Chemistry, Oregon State University, Corvallis. Oregon 97331, and Department of Chemistry, Institute of Mathematical and Physical Sciences, University of Tromsca, N-9001 Tromso, Norway (Received: December 29, 1986)
Core-ionization energies and Auger kinetic energies of fluorine, chlorine, bromine, and iodine in haloethane and haloethene have been used to study the factors that affect the relative ability of ethane and ethene to accept charge at a substituent site. The difference in this ability between the two types of molecules is due almost entirely to the difference in the initial-state charge distribution and almost not at all to differences in the valence-electron rearrangement that accompanies addition of charge to the substituent (in this case, by core ionization). Ab initio calculationsgive results that agree with the experimentally measured quantities and with the interpretation. In the initial state the charge on the halogen is calculated to be less negative in haloethene than in haloethane; both A and u electrons are involved in this charge difference. The contribution to final-state relaxation in haloethene from polarization of the carbon-arbon T bond is approximately matched in haloethane by polarization of the additional carbon-hydrogen u bond in the CY position. Hence, the relaxation energies are nearly the same in the two molecules.
Introduction Such fundamental chemical properties as acidity, basicity, ionization energy, strength of hydrogen bonding, and rates of acidand base-catalyzed reactions depend on the ablity of a molecule to accept charge at a particular site. Traditionally, these properties have been understood in terms of field and resonance effects, which can, in turn, be related to various u parameters.' During the past 15 years, however, emphasis has shifted to the role played by initial-state charge distribution and final-state charge rearrangement in determining these properties.2-8 The effects of charge distribution and rearrangement can be measured experimentally by comparison of core-ionization energies with either Auger kinetic energies6-'s9 or gas-phase Such techniques have been used recently to measure the influence of initial-state charge distribution and final-state charge rear+ Oregon State University. 'University of Tromse.
0022-3654/87/2091-3969$01 .50/0
rangement on the relative electron-donating ability of aliphatic and aromatic rings' and on the relative acidities of organic acids (1) (a) Johnson, C. D. The Hammett Equation; Cambridge University Press: Cambridge, 1973. (b) J a m , H. H. Chem. Rev. 1953, 53, 191. (2) (a) Brauman, J. I.; Blair, L. K. J. Am. Chem. SOC.1970, 92, 5986. (b) Brauman, J. I.; Blair, L. K. J . Am. Chem. SOC.1971.93, 3911. (c) Brauman, J. I.; Riveros, J. M.; Blair, L. K. J . Am. Chem. SOC.1971, 93, 3914. (d) Yamdagni, R.; Kebarle, P. J . Am. Chem. SOC.1973,95,4050. (e) Hiraoka, K.; Yamdagni, R.; Kebarle, P. J . Am. Chem. SOC.1973, 95, 6833. (3) (a) Kollman, P. A.; Allen, L. C. Theor. Chim. Acta 1970, 18, 399. (b) Morokuma, K. J . Chem. Phys. 1971, 55, 1236. ( c ) Dreyfus, M.; Pullman, A. Theor. Chim. Acta 1970,19, 20. (d) Kollman, P. A. Modern Theoretical Chemistry; Schaefer, H. F., 111, Ed.; Plenum: New York, 1977; Vol. 4, pp 109-151. (e) Davis, D. W.; Singh, U. C.; Kollman, P. A. J . Mol. Struct. 1983, 10.5, 99. (f) Davis, D. W. J . Mol. Struct. 1985, 127, 337. (4) (a) Martin, R. L.; Shirley, D. A. J . Am. Chem. SOC.1974, 96, 5299. (b) Davis, D. W.; Rabalais, J. W. J . Am. Chem. SOC.1974, 96, 5305. (c) Davis, D. W.; Shirley, D. A. J . Am. Chem. SOC.1976, 98, 7898. (d) Davis, D. W.; Shirley, D. A. J . Electron Spectrosc. Relat. Phenom. 1974, 3, 137. (5) Smith, S. R.; Thomas, T. D. J . Am. Chem. SOC.1978, 100, 5459.
0 1987 American Chemical Society
3970 The Journal of Physical Chemistry, Vol. 91, No. 15, 1987
Siggel et al.
TABLE I: Experimental Halogen Core-IonizationEnergies ( I ) , Auger Kinetic Energies (K),Relative Initial-State Potential Energies (A V ) , and Relative Final-State Relaxation Energies (AR)O molecule Ib KC AId A K ~ Avd A R ~ fluoroethene fluoroethane
693.28 692.28
599.35 600.51
1 .oo
-1.16
0.92
-0.08
chloroethene chloroethane
206.48 205.91
2374.81 2375.49
0.57
-0.68
0.52
-0.05
bromoethene bromoethane
189.90 189.41
1377.70 1378.19
0.49
-0.49
0.49
iodoethene iodoethane
626.76 626.40
506.41 506.84
0.36
-0.43
0.32
0.00 -0.04
"All values in eV. 6Fluorine Is, chlorine 2p3/,, bromine 3p3/2, iodine 3d5,2. CFluorineKL,L,, chlorine KL2,3L2,3ID,bromine L3M4,,M4,,'G, iodine M4N,,,N4,,'G.
Relative to corresponding saturated molecule.
and alcohols.* Although the total acidity and electron-donating power depend on a combination of initial- and final-state effects, the differences between aliphatic and aromatic rings and between acids and alcohols are due almost entirely to the initial-state charge distribution. The effect of polarization of the a electrons in the compounds with double bonds is approximately offset by additional polarization in the saturated compounds arising from the greater number of electrons (associated with the extra hydrogens needed to saturate the double bonds). In the initial state there is backdonation from the a orbitals of electronegative substituents to the a systems of the aromatic rings. Since a back-donation is not possible for the aliphatic compounds, the substituents are less negative when attached to an aromatic ring than when attached to an aliphatic ring. Similar behavior is seen in the comparison of the acidity of phenol with that of cyclohexanol.* Here we report an extension of these studies to a comparison of ethane and ethene, as representative of saturated and unsaturated aliphatic hydrocarbons. The experimental quantities that provide the information of interest are the core-ionization energies and Auger kinetic energies for the halogens in haloethane and haloethene. As in the previous ~ t u d i e s , ~the . ~ halogen ,~ provides a probe of the response of the rest of the molecule to changes in charge at a substituent site. We have compared the experimental ionization energy shifts, initial-state potentials, and final-state relaxation energies with those calculated from ab initio theory. There is good agreement between the experimental and theoretical results. An analysis of the theoretical calculations provides insight into the factors that influence the ability of these molecules to accept charge at a substituent site. We find that the extra relaxation energy associated with the a bond in haloethene is matched by a nearly equal relaxation energy associated with the extra carbon-hydrogen bond in the a position on haloethane. For the initial state, the calculations show that both the a and 0 charges are less negative in haloethene than in haloethane.I0
Experimental Section The experimental procedures and methods of data analysis are essentially identical with those previously d e ~ c r i b e d .Samples ~ excited with A1 and Mg K a X-rays were calibrated by using the Ne 1s and 2s lines.'1a For samples excited with Ag La, radiation (2984.34 calibration was based on Ne KLL('D) and 1s lines.'Ia Some of the Auger spectra were also obtained by using ~
(6) Aitken, E. J.; Bahl, M. K.; Bomben, K. D.; Gimzewski, J. K.; Nolan, G. S.; Thomas, T. D. J . Am. Chem. SOC.1980, 102, 4873. (7) Nolan, G. S.; Saethre, L. J.; Siggel, M. R.; Thomas, T. D.; Ungier, L. J . Am. Chem. SOC.1985, 107, 6463. (8) Siggel, M. R.; Thomas, T. D. J . Am. Chem. SOC.1986, 108, 4360. (9) Saethre, L. J.; Thomas, T. D.; Gropen, 0. J . Am. Chem. SOC.1985, 107, 2581. (IO) Although one does not usually consider haloethane to have r orbitals, it is convenient to separate the orbitals into those in the CCX plane, which we will refer to as (I orbitals, and those with a node in the CCX plane, which we will refer to as r orbitals. (1 1) (a) Saethre, L. J.; Thomas, T. D.; Ungier, L. J . Electron Spectrosc. Refat.Phenom. 1984, 33, 381. (b) Keski-Rahkonen, 0.;Krause, M . 0. J . Electron Spectrosc. Relat. Phenom. 1916, 9, 371.
electron impact rather than X-ray excitation. In these cases calibration was based on simultaneous measurement of the position of Ne KLL (ID)and Ar LMM (ID) lines excited from neon and argon that were mixed with the sample of interest.'Ia On the basis of estimated standard deviations and reproducibility of consecutive runs, we estimate an overall uncertainty of about 0.05 eV.
Experimental Results and Discussion The experimentally determined halogen core-ionization energies, I, and Auger kinetic energies, K, are listed in Table I. Also shown here are the values for the haloethenes relative to those for the haloethanes (AI and AK). The analysis of these results is based on the fact that the effect of final-state rearrangement on the Auger energies is greater than its effect on core-ionization energies.I2 As has been shown elsewhere,'* the shift in core-ionization energy for the same atom in two different molecules is given by the expression AI = A V - A R
(1)
and the shift in Auger kinetic energy by
AK = -AV
+ 3AR
(2)
In these expressions V represents the effect of an initial-state potential, V / e , on the ionization energy at the site of the core electron and R the contribution to the ionization energy by electron rearrangement that accompanies core-electron removal. With these equations the values of AV and AR are readily obtained from the experimental quantities AI and AK. These are also listed in Table I. Inspection of the last column of Table I shows that the halogen relaxation energies, AR,for the unsaturated compounds relative to the saturated ones are essentially equal to zero. In the ring compounds previously disc~ssed,~ this effect was attributed to the near equality of the extra a polarizability in the aromatic compounds and the additional polarizability of the extra hydrogens (and their electrons) in the aliphatic ones. We will see in the discussion of the theoretical results that the same explanation applies here. The next-to-last column of Table I shows that AV is positive in each case, implying a more positive potential at the halogen atom in the unsaturated species than in the saturated ones. In the ring compounds we showed that this is due to back-donation of a electrons from halogen to the a system of the unsaturated m ~ l e c u l e .A~ similar effect is present here, but we will also see that the halogen m orbitals are less occupied in haloethene than in haloethane. Thus, both u and a orbitals contribute to a less negative halogen in haloethene. The monotonic decrease in AV from the fluoro to the iodo compounds does not, however, mean that the charge difference decreases as we move down the halogens. Because of the increasing size of the halogens as we go down the periodic table, the potential energy Vis less affected by a given (12) (a) Shirley, D. A. Phys. Reu. A 1973, 7 , 1520. (b) Wagner, C. D. Faraday Discuss. Chem. Soc. 1975,60, 291. (c) Siegbahn, H.; Goscinski, 0. Phys. Scr. 1976, 13, 225. (d) Thomas, T. D. J . Electron Spectrosc. Relat. Phenom. 1980, 20, 117.
Electron-Donating Ability of Ethyl and Ethenyl Groups
The Journal of Physical Chemistry, Vol. 91, No. 15, 1987 3971
orbital
[(calcdl
V(calcdl
R(ca1cd)
fluoroethene fluoroethane
molecule
F Is F 1s
694.38 693.37
716.32 715.13
21.94 21.76
chloroethene chloroethane
C1 2p CI 2p
206.84 206.24
218.21 217.58
11.37 11.34
bromoethene
Br 3d Br 3p Br 3d Br 3p
76.43 193.12 75.90 192.56
86.17 202.27 85.60 201.70
9.74 9.15 9.70 9.14
13d 13d
642.34 641.96
660.75 660.40
18.41 18.44
bromoethane iodoethene iodoethane
'All values in eV.
I
I
TABLE II: Results from ab Initio Calculations of Core-Ionization Energies ( I ) , Ground-State Potential Energies ( V), and Relaxation Energies (R)'
I
a
u 208
207
206
Iexpt
I
I
I
I
I
209
lev) I
I
,
I
I
I
charge transfer for iodine than for fluorine. A simple charge analysis, described in a subsequent section, shows that the charge difference is nearly constant as we move from the fluoro to the iodo compounds.
Theoretical Results and Discussion U p to now, only a few comparisons have been made between the values of AV and AZ? derived from measurements of Auger and core-ionization energies and the corresponding values from electronic structure theory. Aitken et aL6showed for an extensive series of chlorine-containing compounds that the experimental values are in reasonable agreement with values obtained from C N D O calculations. They also found that relaxation energies determined experimentally for HCl and ClF (relative to CH,Cl) are in agreement with those calculated by ab initio theory.13 More recently, Saethre, Thomas, and Gropeng have shown that ab initio calculations of AI,AV, and AR for HX relative to Xzgive results that are in agreement with the values obtained from experiment. In view of the small number of such comparisons that exist and the fact that they have been made for relatively simple molecules, we have done a b initio calculations of AZ, AV, and AR for the molecules discussed here. Values of I are obtained by taking the difference between the calculated energies of the core-ionized molecule and the molecule in its ground state, that is, by the ASCF method. Ionization energy shifts, AI, between different compounds are readily calculated from these. Values of AV are taken to be the differences in orbital energies for the appropriate ground-state molecules. Relaxation energies may then be calculated from eq 1, AZ? = A V - AZ. Ab initio LCAO-SCF calculations for these molecules were carried out using the Hartree-Fock method with the MOLECULE-ALCHEMY package.I4 The atomic orbital basis set for hydrogen is 3s (ref 15); for carbon and fluorine, 7s3p (ref 16); for chlorine, 10s6p (ref 16); for bromine, 18s8p5d (ref 17); and for iodine 25sllp7d (ref 18). These sets were contracted to double-r quality. For the chloro compounds, calculations were done both with and without p and d polarization functions (p on hydrogen atoms and d on all other atoms). Since the results obtained with these additional functions agree better with experiment than those without, we have included the polarization functions in the calculations for all of the other halogen-containing molecules. The exponents for the polarization functions for hydrogen and the halogens are the same as in ref 9. The exponents for carbon were energy optimized for ethene and ethane, and the (13) Adams, D. B. J . Electron Spectrosc. Relat. Phenom. 1911, 10, 247. ( 14) The MOLECULE-ALCHEMY program package incorporates the MOLECULE integrals program written by J. Almldf and the ALCHEMY programs written by P. Bagus, B. Lin, M. Yoshimine, D. MacLean, and modified by P. Bagus and U. Wahlgren. (15) Huizinaga, S . J . Chem. Phys. 1965, 42, 1293. (16) Roos, B.; Siegbahn, P. Theor. Chim. Acta 1970, 17, 209. (17) Gropen, O., unpublished results. (18) Stramberg, A.; Gropen, 0.; Wahlgren, U. J . Comput. Chem. 1983, 4, 181.
I
I
1.0 2.0 AV (eV1 Figure 1. Comparison of experimental and theoretical quantities for chlorine. (a) Differences between calculated and measured chlorine 2p ionization energies plotted against experimental values: closed circles, calculated with polarization functions; crosses calculated without polarization functions. Lines represent least-squares fits to the two sets of results. (b) AR vs. A E closed circles, experimental results; crosses, theoretical results.
0.0
mean value, 0.90, was used for carbon in all calculations. Results of these calculations are given in Table 11. Halogen Zonization Energies. We consider first the results for the chlorine-containing molecules, for which we have the most extensive set of measurements and calculations. A comparison of theoretical and experimental ionization energies is shown in Figure l a , where we have plotted the difference between the theoretical and experimental ionization energies vs. the experimental values for the chlorine 2p3jZlevel. We have included in this figure results that have been previously reported6 for HCl, Clz, ClF, as well as a value for ICl.19 Two sets of results are shown, the lower set being those obtained with polarization functions and the upper those obtained without. The first are within 0.3-0.5 eV of the measured values. This close agreement is certainly fortuitous since the calculations do not take into account spin-orbit splitting (or other relativistic effects) or electron-electron correlation. The lines drawn through the two sets of points represent least-squares fits. Values for CIF,Zo are also included in the least-squares calculations but are not shown in the figure because of the very high ionization energy (213.02 eV). The point calculated without polarization functions lies 0.07 eV above the upper line, but that calculated with polarization functions falls right on the lower solid line. The slope for the lower line of 0.10 indicates that the theoretical shifts in ionization energy are 10% greater than the experimental ones. A similar effect has been reported for carbonz1and sulfur.zz Adams and Clarkz1found that the agreement between the theoretical and experimental shifts is improved by using a more flexible basis set. The deviation of the points from the lower least-squares line in Figure l a is quite small (root mean square = 0.02 eV) and is comparable to the (19) Jolly, W. L.; Bomben, K. D.; Eyermann, C. J. At. Data Nucl. Data Tables 1984, 31, 433. (20) Carroll, T. X.;Thomas, T. D., unpublished results. (21) Adams, D. B.; Clark, D. T. Theor. Chim. Acta 1973, 31, 171. (22) Theodorakopoulos,G.; Petsalakis, I. D.; Csizmadia, I . G.; Robb, M. A. J . Mol. Struct. 1984, 110, 381.
3972
The Journal of Physical Chemistry, Vol. 91, No. 15, 1987
Siggel et al.
TABLE III: Comparison of Relative Experimental and Theoretical Halogen Core-IonizationEnergies ( A I ) , Ground-State Potential Energies (AV), and Relaxation Energies (AR)" molecule Al(expt1) Al(calcd) AV(expt1) AV(calcd) AR(expt1) AR(calcd) fluoroethene 1.oo 1.01 0.92 1.19 -0.08 0.18
fluoroethane chloroethene chloroethane
0.57
0.60
0.52
0.63
bromoethene bromoethane
0.49
0.56
0.49
0.57
iodoethene
0.36
0.38
0.32
0.35
-0.05
0.03
0.00
0.01
-0.04
-0.03
iodoethane
"All values in eV
uncertainties of the measurements. Thus, the theory accounts quite well for the variation in ionization energy for compounds in which chlorine changes from an electron-withdrawing ligand to an electron-donating central atom. The upper data, calculated without polarization functions, have a root-mean-square deviation from the least-squares fit (dashed line) of 0.09 eV, which is considerably worse than for the results obtained with polarization functions. Rather than show similar plots comparing experimental and theoretical values of AVand AR,we have found it instructive to plot AR vs. AV. These two quantities depend on rather different characteristics of the molecules and are not well correlated with one another. It is important to demonstrate that theory and experiment agree on the irregular relationships between these two quantities. The available experimental values of AR are plotted against those for AV (both relative to HCl) as the solid points in Figure lb. They are joined, in the order of increasing AV, by a solid line, which serves to emphasize the erratic relationship between the two quantities. The theoretical results (with polarization functions) are shown as crosses and connected by a dashed line joining points in the same order as the experimental ones. We see not only that the overall pattern is reproduced but that the theoretical and experimental quantities agree well with one another. This close agreement indicates that the experimental quantities AVand AR do indeed have the physical significance that we have attached to them. Results for bromine and iodine-containing compounds are quite similar to those for chlorine. The theoretical calculations slightly overestimate the values of AZ, but the experimental relationships between AR and AV are reproduced well by the theoretical calculations. A comparison of experiment and theory for the fluorine-containing compounds is shown in Figure 2. In Figure 2a, we see that the theoretical 1s ionization energies are about 1 eV higher than the experimental ones. Figure 2b shows that experiment and theory approximately agree on the overall relationships between AVand AR,although the agreement is not so goad as for the other halogens. As we have noted e l ~ e w h e r e the , ~ range of AI? for fluorine is large, and second-order corrections that affect the experimental values of AV and AR may be necessary. We see from Figures 1 and 2 and from the data given in Table I1 that the calculated values of V for haloethenes are always positive relative to those for haloethanes and that R is nearly the same for the two types of compounds. These results are summarized in Table 111, where we see that there is close agreement between the experimental and theoretical values of AZ, AV, and AR. Analysis of Relative Relaxation Energies. As we have mentioned above, we believe that the similarity of relaxation energies in the two types of compounds arises from a match between the polarizability of the 7r electrons in the doubly bonded systems and the contribution from the extra electrons present in the singly bonded systems. Our theoretical calculations for haloethane and haloethene support this conclusion. Table IV shows the change in gross electron population for these molecules when a core electron is removed from the halogen. Although the gross electron populations do not necessarily provide
692
693
694 L D t
695 (eV)
696
691
7 0
1
2 3 AV (eV)
4
5
6
Figure 2. Comparison of experimental and theoretical quantities for fluorine. (a) Difference between calculated and measured fluorine 1s ionization energies plotted against experimental values. (b) AI? vs. AV. Closed circles represent experimental results, and crosses represent theoretical results. TABLE I V Calculated Change in Electron Population following Core Ionization molecule halogen Cf1P C(2Y H(l)b H12Y
CHzCHF CH3CHZF
0.39 0.38
0.19 0.18
-0.18 0.04
-0.16 -0.32
-0.24
CHZCHCI CHBCHZCl
0.37 0.37
0.06 0.03
-0.12
0.06
-0.12 -0.22
-0.20 -0.23
CH,CHBr CH3CH2Br
0.39 0.39
0.02 -0.02
-0.11
-0.11
0.05
-0.20
-0.19 -0.22
CHZCHI
0.34 0.35
-0.01
-0.08 0.05
-0.09 -0.16
-0.16 -0.19
CH3CHJ
-0.05
-0.28
"C(1) is the carbon to which the halogen is attached. bSum of all hydrogens attached to C(1). 'Sum of all hydrogens attached to C(2). an accurate picture of the charge distribution in the molecule, the changes in population may give a reasonable qualitative picture of the charge flow that accompanies core ionization. To simplify the presentation, we have combined the charge flow for all the hydrogens attached to a given carbon into a single number. Several trends can be seen from inspection of these results. First, there is virtually no difference in the charge transfer to the halogen for corresponding alkanes and alkenes. Since this charge transfer is the major contribution to the relaxation energy, this
The Journal of Physical Chemistry, Vol. 91, No. 15, 1987 3973
Electron-Donating Ability of Ethyl and Ethenyl Groups
derived from experiment agree reasonably with those from theory for fluorine and chlorine, this is not the case for bromine and iodine. The charges are sensitive to the choice of k, especially for iodine, for which the difference between k and e2/RcI is small.
TABLE V Calculated Halogen Charge Differences" halogen fluorine chlorine bromine iodine
Aq,(calcd)
AqJcalcd)
Aq,,,(ca!cd)
Aq(expt!)
0.034
0.013
0.047
0.031 0.028
0.022 0.038
0.053
0.04 0.04
0.066
0.04
0.019
0.075
0.094
0.03
"Aq = q(ha1oethene) - q(ha1oethane). result is consistent with the observation that AR is nearly zero for all pairs of molecules. Second, we see that the changes on C(l) and on H(2) are nearly the same for alkanes as for alkenes. The major differences between the two systems are found in the behavior of electrons on either C(2) or H ( 1). Inspection of the data for C(2) shows significant loss of electrons from this position in the alkenes, which is not present in the alkanes. This charge transfer results from polarization of the a electrons in the double bond. On the other hand, the results for H ( l ) show twice as much negative charge lost by the two hydrogens in the alkanes as by the single hydrogen in the alkenes. In this case we have polarization of the carbon-hydrogen u bond. Thus, the relaxation associated with the double bond in the one case is essentially the same as that associated with the extra hydrogen attached to C( 1) in the other. The relevant factor seems to be the number of bonds rather than whether they are u or a. Analysis of Initial-State Charge Distributions. In our earlier analysis7 we concluded that the positive values of AYfor aromatic relative to aliphatic rings arise because of back-donation of a electrons from the electronegative substituent to the 7 system of the aromatic ring. We now consider the same problem for haloethane and haloethene by examining the difference in halogen charge in the two kinds of molecules. The calculated differences are given in Table V, where we show the a and u differences as well as the total difference.'O Inspecting the next-to-last column of Table V, we see that the total charge on the halogen is always more positive in haloethene than in haloethane. The more positive value of V in haloethene arises principally from this charge difference. From the first two columns we see that only for fluorine and chlorine is the difference in a charge the principal source of the extra positive charge. In all cases, the u charge is less negative in haloethene than in haloethane, with the difference being most pronounced for the iodo compounds. This effect may result from the higher electronegativity of the carbon sp2 bond in haloethene compared to the sp3 bond in haloethane. We use a simplified point-charge model7 to estimate charges on the halogens from the experimental values of AY (in this case, relative to the diatomic halogen). W e assume that the entire ground-state halogen charge is withdrawn only from its nearest n e i g h b ~ r .The ~ halogen charge, q, can then be calculated from the expression q = AV/(k - e 2 / R c x )
(3)
where k is the change in potential at the halogen core when a valence electron is removed and Rcx is the distance between the halogen and the adjacent carbon. Although this model ignores details of the charge distribution, it contains the essential features. The constant k is often equated to ( l / r ) ,where r is the valence radius. For our calculations we have used values from C a r l ~ o n . ~ ~ We are concerned here only with the differences between the halogen charge in haloethene and that in haloethane. These are given in the last column of Table V. We see that the halogen in haloethene is always less negative than in haloethane. However, the charge variation with halogen type is much less pronounced than the variation of AY(Tab1e I) because the values of k decrease as we go down the halogen series. Although the charge differences
(23) Carlson, T. A. Photoelectron and Auger Spectroscopy; Plenum: New York, 1975; pp 188-193.
ConcIusion The results presented here as well as other related ~ o r k lead ~ - ~ to two important conclusions. First, ab initio calculations give firm support to the idea that we can obtain values of AVand AR from combining measured core-ionization energies with either Auger energies9 or gas-phase acidities.8 Second, experiment and theory agree that the relaxation energy contribution is approximately the same for compounds with double bonds as for compounds of similar structure with no double bonds and that the compounds with double bonds are less electron donating than those witho~t.~,~ Considering the first of these points, we note that ab initio theory at the level we are using it gives predicted core-ionization energy shifts in agreement with the experimental values. This agreement gives us confidence that we are using theory that adequately describes the principal processes taking place when charge is removed from or added to the molecule. The theoretical results are easily resolved into initial- and final-state contributions, since these are calculated separately. We find that these agree well with those derived from the experimental quantities using the model that has been developed for this purpose.12 This agreement gives us confidence that the experimental values of AV and AR do indeed have the physical significance that has been assigned to them. Turning to the second point, we note that these conclusions have significance beyond the interpretation of core-ionization energies. These energies correlate well with other chemical properties, such as acidityssz4and b a s i ~ i t y . ~ "The ~ , ~factors ~ that influence these more familiar properties also influence core-ionization energies. The conclusion that there is essentially no difference in relaxation energy between compounds with and without double bonds applies not only to the comparisons of core-ionization energies considered here and in our previous paper7 but also to the comparison of the acidities of phenol and cyclohexanol or acetic acid and 2-propan01.~ Similarly, that compounds with double bonds are less electron donating than those without is the principal factor in determining not only the relative ionization energies in haloethane/haloethene or halobenzene/halocyclohexane7 but also the relative acidities of organic acids and organic alcohols.6 To us, the most surprising result from these studies, and the most difficult to accept, is that molecules with double bonds are no more polarizable than those without. Our prejudice that a bonds should be more polarizable than (I bonds has been shared by many people with whom we have discussed the question. The expeiimental and theoretical results presented here and elsewhere7~* are unambiguous in this respect. In the ring compounds7 we have seen that the contribution from polarization of the a electrons in the aromatic compounds is matched by contributions from the extra hydrogens in the aliphatic ones. In the compounds considered here, the a polarization in ethene is matched in ethane by an equal contribution from the extra hydrogen attached to the CY carbon. Acknowledgment. This work has been supported in part by the U S . National Science Foundation and the Norwegian Council for Science and Humanities. Grants from NATO (RG 103/84) and the Norwegian Council for Scientific and Industrial Research, which have allowed the authors to consult face to face on several (24) Jen, J. S.; Thomas, T. D. J . Am. Chem. SOC.1975, 97, 1265. (25) (a) Carroll, T.X.; Smith, S . R.; Thomas, T. D. J . Am. Chem. SOC. 1975, 97, 659. (b) Mills, B. E.; Martin, R. L.; Shirley, D. A. J . Am. Chem. SOC.1976, 98, 2380. (c) Benoit, F. M.; Harrison, A. G. J . Am. Chem. SOC. 1977, 99, 3980. (d) Cavell, R. G.; Allison, D. A. J . Am. Chem. SOC.1977, 99, 4203. (e) Ashe, A. J., 111; Bahl, M. K.; Bomben, K. D.; Chan, W.-T.; Gimzewski, J. K.; Sitton, P. G.;Thomas, T. D. J . Am. Chem. SOC.1979,101, 1764. (0 Brown, R. S.; Tse, A. J . Am. Chem. SOC.1980, 102, 5222. (g) McMahon, T. B.; Kebarle, P. J . Am. Chem. SOC.1985, 107, 2612.
J . Phys. Chem. 1987, 91, 3974-3977
3974
occasions, are gratefully acknowledged. We thank Stephen E. Anderson for completing the last few experimental measurements. T.D.T. thanks the University of Liverpool for providing its hospitality and the Royal Society for its support while this article
was being written. Registry No. CH3CH2F,353-36-6; CH3CH2CI,75-00-3;CH3CH,Br, 74-96-4; CH3CH21,75-03-6; fluoroethene, 75-02-5; chloroethene, 7501-4; bromoethene, 593-60-2;iodoethene, 593-66-8.
Solvent Effect on the Intramolecular Hydrogen Bond Strength and on the Isotopic Ratio v,,+/v,,+ in a Trisubstituted Mannich Base Maria Rospenkt and Th. Zeegers-Huyskens* Department of Chemistry, University of Leuven. Celestijnenlaan 200F, B- 3030 Heverlee, Belgium (Received: January 20, 1987)
The electronic (300-440 nm) and the infrared spectra (3000-700 cm-') of 2-(N,N-dodecylaminomethyl)-3,6-dichloro-4nitrophenol and its deuteriated OD analogue have been studied in 10 different organic solvents. The strength of the intramolecular NH+-O- hydrogen bond decreases with the Onsager parameter of the solvent. The v ~ ~ + . values . . ~ - ranging from 2800 to 2700 cm-' suggest that the hydrogen bond is of medium strength. The isotopic ratio VNH+/VND+ values varying from 1.322 (CC14) to 1.350(C,H4Cl,) suggest a double-minimum curve with a relatively high barrier for the proton motion. The experimental - ~ - are mprkedly higher for the intra- than for the intermolecular results show that, for the same isotopic ratio, the v ~ ~ +values hydrogen bonds. From literature data, the same trend can be found for intra- and intermolecular OH-0 bonds. This effect can be accounted for, at least qualitatively, by the nonlinearity of the intramolecular hydrogen bonds.
Introduction As shown by recent infrared data,'-4 dipole moment measurements, and X-ray ortho Mannich bases are classes of compounds characterized by stable intramolecular hydrogen bonds. By introducing
is of the proton-transfer type. The isotopic ratio is intimitely related to the strength of the hydrogen bond, to the intermolecular distances, and to the shape of the potential curve for the proton motion. It seemed therefore interesting to investigate the solvent effect on the isotopic ratio vNH+/VND+ and to compare the results with those previously reported for normal OH-N bonds; for these, the isotopic ratio vOH/VOD has been shown to decrease with the polarity of the m e d i ~ m . ~
Experimental Section The ultraviolet spectra were taken with a Cary 219 spectrophotometer using I-cm cells and concentrations between 2 X low3 various X groups into the phenyl ring and changing the amines and 8 X M. The infrared spectra were recorded with a condensed with the appropriate phenols, it is possible to modulate Perkin-Elmer 580B spectrophotometer using cells with KBr within a broad scope the acidic and basic properties of the inwindows of 0.05-0.2 cm. In the weakly polar solvents, the solteracting centers. According to ultraviolet data, the intramolecular ubility of DCNMBHp, is very low and maximal concentrations proton-transfer constant KpTdepends on the substituents implanted on the phenolic ring, on the solvent, and on the t e m p e r a t ~ r e . ~ - ~ (about 0.015 M) were used in order to obtain representative spectra. When three or four chlorine atoms are implanted on the aromatic DCNMBH was synthesized recently in the laboratory of Proring, the hydrogen bond is predominantly of the OH-N type fessor L. Sobczyk by a described procedure." The deuteriated whatever the polarity of the ~ o l v e n t . ~In , ~ 2-(N,N-dimethylanalogue was obtained by dissolving the base in C H 3 0 D under aminomethyl)-4-nitrophenol,two NH+.-O- bonds related through nitrogen atmosphere. This operation was repeated several times an inversion center form a cyclic dimer in the solid state while and about 7 0 4 0 % deuteriation was achieved. Owing to the rapid in weakly polar solvents, intramolecular OH-N hydrogen bonding exchange with the hydrogen atoms, the infrared spectra of is predominating.* The purpose of this work is to investigate the solvent effect on the strength of the hydrogen bond and in 2-(N,N-dodecyl(1) Sucharda-Sobczyk,A.; Sobcyzyk, L. Bull. Acad. Pol. Sci., Ser. Sci. aminomethyl)-3,6-dichloro-4-nitrophenoland its O D deuteriated Chim. 1978, 26, 549. analogue: ( 2 ) Koll, A.; Glowiak, T. J . Crystallogr. Spectrosc. Res. 1985, IS, 411. (3) Rospenk, M.; Zeegers-Huyskens, Th. Bull. SOC.Chim. Belg. 1985, 94, 469
