NOTES
1927
Ionic and Partial Double-Bond Character of Carbon-Chlorine Bond in
$
a -141
1,l -Difluorovinyl Chloride h I
g
L i -18
by Suresh Chandra
e:
Physks Department, Allahbad University, Allahbad, India Accepted and Transmitted by The Faraday Sosociety
5
-26 -26
- 22
-IS
- 14
- 10
alp chemical shift of R P O F , ppm.
Figure 1. A plot of the S I P chemical shifts of variously substituted phosphines us. the chemical shifts of quaternary triphenylphosphonium compounds having the same substituent. The dotted line corresponds to the theoretical slope (see text).
If the P-0 ?r bonding and the OPO bond angles in the phosphonates were not affected by changing the R substituent, a plot of 31Pchemical shifts of the phosphonate anions us. those of the corresponding phosphonium cations, RPR'3+, should give a straight line However, as shown in Figure 1, this is of slope far from being the case, so that it appears that varying the R organic substituent does affect the P-0 bond order to a small extent. In other words, the great sensitivity of the 31Pchemical shift to the occupation of the phosphorus d, orbitals renders the phosphonates a poor tool for estimating the electron-donating ability of organic groups. Both bond-distance datas and the observed3 31Pchemical shifts of triphenylphosphine and the tetraphenylphosphonium ion indicate that there is little or no ?r character in the P-C bond between a phenyl group and a phosphorus atom. Therefore it seems reasonable to ascribe essentially all such ?r effects to the P-0 bonds of the phosphonate anions. This is borne out by the rather good agreement of the A6 valued between the various quaternary triphenyl- and trialkylphosphonium cations, shown in Table 11. If there were much a bonding between the phenyl group and the phosphorus and if changing the R substituent affected this ?r bonding, these values of AS would be quite inconsistent. (8) In the compound (C&) (CHa)(S)PP(S) (CHI) (CaHs), the P-C distances for the methyl groups are 1.82 A whereas they are 1.88 A for the phenyl groups, according to P. J. Wheatley, J . Chem. Soc., 523 (1960). r bonding should shorten the bond. Other X-ray work has shown the P-C distance in (CHs)aP to be 1.84 A and in (CsHa)aP to be 1.83 A.
(June 6 , 1966)
The microwave spectrum of 1,l-difluorovinyl chloride has been studied by Jenkins and Sugdenl and a structure for this molecule was given. A new and more probable structure for 1,1-difluorovinyl chloride has been proposed in order to get better and more probable molecular characteristics. From this structure a new value of the angle 8 between the a axis and the G C 1 bond has been found. The principal quadrupole coupling constants, ionic and partial double-bond character of the carbon-chlorine bond, and percentage of s character of the chlorine bond orbital have also been evaluated. The bond lengths and bond angles are intimately related to their bond environment. For a given bond environment they are remarkably constant in different molecules but vary systematically with the changes in the bond environment. However, small changes are observed when C1, Br, and I atoms are adjacent to the bond.2 So, for purposes of calculating approximate moments of inertia, structural parameters for 1,ldifluorovinyl chloride (CFFCHCI) are deduced from those molecules which have very nearly the same structure as C F F C H C l . The carbon-fluorine and carboncarbon bond lengths and the corresponding angles (bond angles) are taken from the structure of the 1,ldifluoroethylene molecule (C2H2F2) and the carbonchlorine and carbon-hydrogen bond lengths and the corresponding bond angles are taken from the structure of the vinyl chloride molecule (CHFCHCI).~ In these calculations, the carbon-carbon bond distance was varied slightly in order to obtain better agreement between calculated and observed moments of inertia. The values of the structural parameters thus obtained (1) D.R. Jenkins and T. M. Sugden, Trans. Faraday Soc., 55, 1473 (1959). (2) A. S. Rajput and S. Chandra, Bull. Chem. Soc. Japan, 39, 1854 (1966). (3) J. Karle and I. L. Karle, J . Chem. Phys., 18, 163 (1950). (4) A. Roberts and W. F. Edgell, ibid., 17, 742 (1950); Phys. Rev., 76, 178 (1950). (5) D.Kivelson, E.B. Wilson, and D. R. Lide, J . Chem. Phys., 32, 205 (1960).
Volume 71, Number 6 May 1967
NOTES
1928
are quoted in Table I. The calculated and observed moments of inertia of CF2=CHW1 and CFpCH"C1 are also compared.
Table 11: Quadrupole Coupling Constants of C F p C H W l f 1 . 3 Mc/sec 18.2 f 1 . 0 Mc/sec zcc = 33.5 f 0 . 8 Mc/sec zEl = -73.94 f 1 . 5 Mc/sec zx3.= 40.44 f 1 . 0 Mc/sec ZYV = 3 3 . 5 f 0 . 8 illc/sec e = 260 io'
zUa= -51.7 Zbb =
Table I : Structural Parameters of 1,l-Difluorovinyl Chloride Bond length, Bond angle
A
C-F := 1.32 C-C = 1.30 C-C1 = 1.726 C-H = 1.08
LFCF = 110" LCCH = 123' 49' LCCCl = 122' 18'
bond, which is very close to the value 6% in vinyl chloride.' It may be mentioned here that eq 1 makes no allowance for the positive charge on chlorine in the structure. The ionic character of a bond AB can be estimated from the electronegativity values (XA and XB) of the two atoms forming the bond. For a two-electron bond, the pair of electrons will divide themselves between the two atoms in the ratio of their relative electronegativities, such that a fraction ~ X A / ( X A XB) of electron atmosphere is on the atom A and a fraction 2xB/ ( X A XB) of electron atmosphere is on the atom B. The ionic character of the bond is then defined as*
Moments of inertia of CFp=CHWl (amu A2) Ib (amu A2) I, (amu A2) Za
Obsd
Calcd
47.20 220.09 267.45
46.93 220.52 267.45
Moments of inertia of CF-CHWl
I. (amu A2)
Ib (amu A*) I, (amu A2)
Obd
Calcd
47.20 226.41 273.76
46.94 226.94 273.88
+
+
ionic character p The characteristic that the a axis passes through the chlorine nucleus was utilized for evaluating the angle e between the a axis and the x axis (along the carbonchlorine bond) which comes out to be 26" 10' from the proposed structure of 1,l-difluorovinyl chloride. If it is assumed that the C-C1 bond coincides with the z principal axis of the quadrupole coupling dyadic, a transformation of axes may be used to find the components of the dyadic in its principal system.' The other two axes must lie in the plane of symmetry (z axis) and perpendicular to it (y axis). The values of the principal quadrupole coupling constants zzz, xvv, and x,, for CFz=CHCWl obtained from the experimental values of xu,, x b p , and xCc1 are listed in Table 11. The value of x,, thus obtained may be compared with that of -70.16 mc/sec for vinyl chloride6 and -74.77 Mc/sec for methyl chlorides and with many other analogous values for chloromethanes and chloroethylenes. Following Goldstein,? an estimate of the partial double-bond character 6 of the carbon-chlorine bond may be made from the relation 6 =
xxz
- x,,
- 3/2eQ~(atomio)
(1)
where -','peQq(atomic) = 164.6 Mc/sec for W1. The result is 4.3% double bond character to the C-CI The Journal of Phyeical Chemistry
=
IX.4 - x B /
(x.4+ XB)
(2)
For polyatomic molecules, the values of group electronegativityg should be used. The group electronegativity value of the radical -CHCF2 is 2.5E1,~ and the electronegativity value of the chlorine atom is 3.00. Using these values, the ionic character of the carbon-chlorine bond in 1,l-difluorovinyl chloride is found to be 8%. The values of the nuclear quadrupole coupling constants may be used for evaluating the hybridization of the atomic orbitals about the nucleus under consideration. The value of the fractional s character of the bonding orbital of chlorine, a2,may be evaluated from the relationlo eQq,, = (1 - a2
+ d2 - P -
8)eQqatomic
(3)
where d2 is the amount of the d hybridization of the bonding orbital of chlorine and e&qatomic is the nuclear quadrupole coupling constant for the free atom. However, the d orbitals make only negligible contributions to the quadrupole coupling constant and its value (6) J. Kraitchman and B. P. Dailey, J . Chem.Phys., 22,1477 (1954) (7) A. Goldstein, ibid., 24, 106 (1956).
(8) J. K.Wilmshurst, ibid., 30, 561 (1959). S. Chahdra, Tetrahedron, 2 2 , 3403 (1966). (10) T.P. Das and E. L. Hahn, "Solid State Physics,'' Supplement 1, "Nuclear Quadrupole Resonance Spectroscopy," Academic Press Inc., New York, N. Y.,1958, p 138. (9) 9. Chandra and
NOTES
iE predicted t o be 5% or less. On the assumption of 5% d character and for the observed values of eQq,, (-73.94 c/sec) and eQqstomic (-109.74 Mc/sec), eq 3 shows that the s character of the chlorine bond orbital is nearly 25%. Aclmowledgment. The author expresses his gratitude and thanks to Professor Krishnaji for his kind supervision of the work.
Experimental Determination of the Electron Affinity of Several Aromatic
Aldehydes and Ketones
1929
of AB, k~ is a pseudo-first-order rate constant representing the loss of electrons due to processes other than attachment to AB, k~ is a pseudo-first-order rate constant representing the loss of AB- due to ail processes other than detachment or dissociation, kl is the rate constant for the attachment of electrons to AB, k-1 is the rate constant for the detachment of electrons from AB-, and K is the electron-capture coefficient. For some compounds, there is one region, designated p, at lower temperatures where k~ >> k-1, so that K = kl/kD
(2)
and K is relatively insensitive to temperature. I n another region, designated a, k-1 >> k~ and, assuming the statistical thermodynamic expression for an ideal gas
by W. E. Wentworth and Edward Chen' Department of Chemistry, University of Houston, Houston, Texas 77004 (Received December 19, 1966)
Recently, Wentworth, Chen, and Lovelock2 have proposed a kinetic model for the processes occurring within the electron-capture detector operated in the pulse-sampling mode and have demonstrated its validity for several aromatic hydrocarbons. Earlier Wentworth and Beckera calculated the electron affinities of some aromatic hydrocarbons using the electroncapture detector data and correlated these electron affinities with their half-wave reduction potentialsSand the energy of their 0-0 tran~ition.~The experimental electron affinities were also compared with several theoretical calculations of the electron affinitie~.~ However, thus far only aromatic hydrocarbons have been studied, although the technique for the calculation of the electron affinities originally given in ref 3 and later modified by the kinetic model2 is applicable to any molecule which forms a stable negative ion. Therefore, the electron-capture detector response to a series of aromatic aldehydes and ketones has been studied as a function of temperature in order to test the validity of the interpretation for a different type of molecule which apparently forms a stable negative ion with respect to electron attachment. For the kinetic model, the following expression can be derived for the response2
where b is the electron concentration before the addition of the test molecule, AB, [e-] is the electron concentration in the presence of AB, a is the initial concentration
A/T'/' exp(EA/kT)kL/kD
(3)
A is a constant which can be calculated from fundamental parameters, EA is the electron affinity, and IC is the Boltzmann constant. In the a region, the electron affinity can be obtained from the slope of a In KT"' vs. 1/T graph.
Experimental Section The procedure and equipment used in this study have been described earlier. The naphthaldehyde-1, benzaldehyde, and acetophenone were Eastman White Label. The phenanthrenealdehyde-9 and naphthaldehyde-2 were obtained from the Aldrich Chemical Co. The solvent used was Eastman's Spectroquality benzene. Results and Discussion The capture coefficients for the compounds studied are plotted in Figure 1 as In KT"' vs. 1/T. All of the compounds except cinnamaldehyde exhibit an a region. Acetophenone and benzaldehyde do not have a /3 region. From the data in the a region, the electron affinities have been calculated from the slopes using an average intercept2v5and a variable intercept. The average value for the intercept from the aromatic hydrocarbons and the compounds used in this study was used and is equal to 14.8. The average intercept (1) This work was used for partial fulfillment of the requirements for the Ph.D. degree at the University of Houston, Houston, Texas. (2) W. E.Wentworth, E. Chen, and J. E. Lovelock, J . Phys. Chem., 70,445 (1966). (3) W. E.Wentworth and R. S. Beoker, J . A m . Chem. SOC.,84,4263 (1962). (4) R. S. Becker and W. E. Wentworth, ibid., 85, 2210 (1963). (5) W. E. Wentworth, W. Hirsch, and E. Chen, J. Phys. Chem., 71, 218 (1967).
Volume 71, Number 6 M a y 1967
