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C O M M U N I C A T I O N S TO THE E D I T O R
Temperature-Dependent Methoxyl an.d Hydroxyl Splitting Constants in the Electron Spin Resonance Spectra of Cation Radiaals'
Sir: The accurate measurement of temperature dependent hyperfine coupling constants in esr spectra can often lead to interesting information on its own account arid may sometimes be explained by physical models. For example, the temperature dependencies of alkali metal splittings give much information about the structure of ion pairs.2 Also, the temperature dependence of alkyl groups may be ascribed to the existence of preferred conformations and may be used to obtain an estimate of their potential barriers to rotat i ~ n . In ~ this communication we wish to report the accurate measurement of the temperature dependencies of the methoxyl and hydroxyl group protons in a number of cation radicals. Reasons for their dependence are suggested and a method for estimating potential barriers to rotation is indicated. The cdtion radicals studied together with their measured temperature dependencies of the methoxyl and hydroxyl proton splitting constants are shown in Table I V 4The radicals were produced in the HZSO4CH3C025and AIC13-CH3N026systems. The spectra were analyzed by previously described methods7 and the splitting constants were found to vary linearly between 200 and 300°K. Two results of particular significance are apparent from a perusal of Table I : first, the temperature dependencies increase along the two series I, 11, I11 and
Table I
Cation radical Hydroquinone (I)
1,4-Dihydroxynaphthalene(11) 9,lO-Dihydroxyanthracene(111) 2,3-Dimethyl-l,$-dihydroxynaphthalene (IV) Durohydroquinone (V) 1,4-Dimethoxybenzene (VI) 1,4-Dimethoxynaphthalene(VII) 9,lO-Dimethoxyanthracene(T'III) Dimethoxydurene (IX)
The Journal of Physical Chemistry
Temperature dependence of the hydroxyl or methoxyl group proton splitting constants, mG OC-1 0.84
1.58
2.87 1.94 1.40 -0.09 0.41 1.16
1.50
VI, VII, VIII. Second, the hydroxyl groups are always more temperature dependent than their methoxyl counterparts (ie., compare I1 and VII). It has been shown previously that I and VI exist as cis and trans isomers a t low temperatures,sp9 indicating a preferred conformation of the hydroxyl and methoxyl groups in the plane of the aromatic ring. The temperature dependence can therefore be considered to arise from torsional oscillations (which change the splitting constants) about a potential minimum, either in or close to the aromatic plane. The amount of temperature dependence will depend upon the depth of the potential well which will in turn depend on two factors. These are the bond order of the >C-0 bond which will tend to constrain the group in the aromatic plane and steric interactions which will try to force the group out of the plane. A combination of these two factors, a decrease in bond order plus an increase in steric interactions, will account for the variations along the series I, 11, I11 and VI, VII, VIII. The generally larger temperature dependence of hydroxyl over methoxyl groups may, to a large extent, be accounted for by the mechanisms of the hyperfine interactions. Thus, when the hydroxyl group is in the aromatic plane the proton splitting arises via a spin polarization mechanism to the unpaired spin on oxygen.'O If the hydroxyl group is twisted out of the plane by some angle e the spin density on oxygen will decrease by a factor cos2 e, thereby decreasing the magnitude of the splitting constant. I n addition the hydroxyl proton may now couple with the unpaired spin on the adjacent carbon atom via a hyperconjugative mechanism. This interaction which will be proportional to sin2 e will be of opposite sign compared to the spin polarization interaction. The two mecha(1) Presented in part at the 157th National Meeting of the American Chemical Society, Minneapolis, Minnesota. Research supported in part by Grant No. NSF-GP-8416 from the National Science Foundation. (2) N. Hirota, J . Phys. Chem., 71, 127 (1967). (3) G. Chapelet-Letourneux, H. Lemaire, R. Lenk, M-A. MarBchal, and A. Rassat, Bull. Chim. SOC.Fr., 3963 (1968). (4) Radicals I V , VII, and VI11 have not previously been reported and will be discussed in greater detail in a later publication. (5) P. D. Sullivan and J. R. Bolton, J . Mag. Res., 1, 356 (1969). (6) W. F. Forbes and P. D. Sullivan, J . Amer. Chem. SOC.,88, 2862 (1966). (7) P. R. Hindle, J. Dos Santos Veiga, and J. R. Bolton., J . Chem. Phys., 48, 4703 (1968). (8) A. B. Barabas, W. F. Forbes, and P. D. Sullivan, Can. J . Chen.. 45, 267 (1967). (9) W. F. Forbes and P. D. Sullivan, Can. J . Chem., 44, 1501 (1966). (10) A. Carrington and H. C. Longuet-Higgins, Mol. Phys., 5 , 448 (1962).
COMMUNICAT~ONS TO THE EDITOR
2791
nisms will combine to give a rapid decrease in the hydroxyl group splitting constant ( ~ 0 ” ) when the group is twisted out of the plane. Thus czOHM = - A (cos20)
+ B (sin2 e)
where A and B are constants and the ( ) brackets denote a time-averaged value. The methoxyl group, on the other hand, couples via a hyperconjugative mechanism to the unpaired spin on oxygen, which is at a maximum when the methoxyl group is in the plane of the ring and will decrease as the methoxyl group is twisted out of the plane such that) aoCHaH =
c (COPe)
The above discussion provides a basis for expecting a greater temperature variation for the hydroxyl group. Further calculations enable us to estimate the depth of the potential well from the temperature dependencies. This may be done by evaluating (cos2 e) for various temperature and barrier heights, using the procedures of Stone and Maki.ll A comparison of calculated and experimental temperature dependencies leads us to estimate the barrier heights for VI, VII, VIII, and I X as 16 i 4,5 f 1, -0.8 and -1.8 kcal/ mol, respectively. Further work on the hydroxyl compounds should enable us to similarly obtain their potential barriers which may then be compared with the potential barriers already obtained by line-width alternation studies.12
Aclcnotdedgments. The author wishes to thank Dr.
J. R . Bolton for his continued encouragement and helpful comments and also Dr. G. Vincow for a copy of his computer program. (11) E. W. Stone arid A. H. Maki, J. Chem. Phys., 37, 1326 (1962). (12) P. D. Sullivan, J . Amer. Chem. SOC.,89, 4294 (1967).
DEPARTMENT OF CHEMISTRY OF MINNESOTA UNIVERSITY MINNEAPOLIS, MINNESOTA55455
PAUL D. SULLIVAN
RECEIVED APRIL10, 1969
The Electron Attachment Cross Section for Hexafluoroacetone
Sir: Among the processes which may occur when an electron suffers a collision with a molecule is electron capture, a negative ion being formed AB
+ e +AB-
Although such interactions have been studied extensively, in few cases has the molecule-ion been suf-
ficiently stable to be detected, the more usual process being the dissociative capture reaction AB
+ e +A- + B
SF6- is the classic example of a stable moleculeion1a2and the SFe- abundance curve has been used to mirror the electron energy distribution and calibrate the electron energy scale. Recently, we reported that hexafluoroacetone formed a stable negative ion3 as a result of secondary electron capture, but we were unable to observe its formation at very low electron energies. Using an improved experimental technique, we now report the formation of the CFsCOCF3- ion at low electron energies following primary electron capture and have obtained a value for the electron attachment cross section of the ketone, relative to that for sulfur hexafluoride. Results were obtained using a Bendix time-of-flight mass spectrometer, Model 3015. The electron energy was measured using a Solatron digital voltmeter LM 1619 and “negative” voltages were obtained by incorporating a 3-V dry cell into the electron energy circuit. Using two channels of the mass spectrometer analog output scanners enabled the SFs- and CF3COCFs- ions to be measured simultaneously on l-mV Kent potentiometric recorders. The electron current was kept constant automatjcally over the energy range studied. Ion source pressures were usually maintained below 5 X 10-6mm. When hexafluoroacetone (HFA) was studied at. electron energies -0 eV a fairly abundant parent ion was observed; admission of sulfur hexafluoride to the ion source considerably reduced the intensity of the parent ion, suggesting that the attachment cross section for the reaction SFa
+e
1 ----f
SFs-
was much greater than for
CFsCOCF3
+ e 2,CF3COCF3-
I n Figure 1 (full circles) we show the data obtained for SFa- ion formation using a 50-50 mixture of HFA and SFs; “negative” voltages were obtained by introducing a 3-V dry cell into the electron energy circuit, The smooth curve reflects the electron energy distribution. I n Figure 1 (open circles) we report our experimental data for the CF3COCF3- ion, the ordinate being 58.9 times more sensitive than that for SFe-. It is apparent that the two curves have a very similar distribution, both reaching a maximum value at the same electron energy. The ketone is slightly broader in the wings; (1) W. M. Hickam and R. E. Fox, J . Chem. Phys., 25, 642 (1956). (2) G. J. Schulz, J . A p p l . Phys., 31, 1134 (1960). (3) LJ. C. J. Thynne, Chem. Commun., 1075 (1968).
Volume 73, Number 8 August 1969
