F. J. SMENTOWSKI AND GERALD R. STEVENSON
340 AGfOZgs.15 of Tm20a(c)to be -428.8 AGfOzgs.lsfor Luz03(c)to be -427.6
f
1.5 kcal/mol and 1.5 kcal/mol.
f
Acknowledgment. The authors are grateful for the partial support of the United States Atomic Energy
Commission and to Professor F. H. Spedding and the Ames Laboratory of Iowa State University for the loan of calorimetric samples of both oxides. They further acknowledge the technical assistance of Dr. John C. Trowbridge in making the measurements.
Temperature-Dependent Electron Spin Resonance Studies.
11.’
Cyclooctatetraene Anion Radical
by F. J, Smentowski and Gerald R. Stevenson Department of Chemistry, Texas A & M University, College Station, Texas 77849 (Received April $6,1968)
The cyclooctatetraene (COT) anion radical (I) has been formed by the alkali metal reduction of COT in liquid ammonia. Ion pairing affects the esr spectra of the anion radical I, influencing the line widths of the individual hyperfine splittings, the activation energy of the line-broadening process, and the spin concentration. The dependence of the hyperfine line width of the cyclooctatetraenide salts is Li+ > Na+ > K+, the reverse order of that found in ether solvents. For the COT-”3-K system, the radical(s) observed are dependent upon the ratio of COT to its dianion (11). Equilibrium constants for the disproportionation (eq 3) in liquid ammonia have been measured. Factors affecting the direction of reaction 3 are considered.
Introduction Ion pairing has been found to play a predominant role in solution Current developments hnvc h e n stimulated hy C S ~ , ” ~conduct,anceJGand opticalGstudics. Much of the esr work on ion-pairing phenomena has recently bcen reviewed.’ Studies8fe of the cyclooctatetraene (COT) anion radical (I) in
I
I1
tetrahydrofuran (THF) and 1,2-dimethoxycthane (DME) have given some insight into ion pairing on these systems. Since details of the preparation of anion radicals in liquid ammonia are now available,’O it was of interest to study the system COT-”3metal. Anion radicals of various organic substrates have been prepared by electrolytic reduction in liquid arnmonia.l1 COT has been reduced electrolytically in liquid ammonia, giving nine lines with aH = 3.28 G.”
Experimental Section Adequate safety precautions must be considered, since the esr sample tubes of these ammonia systems are at approximately 11 atm at 30”. COT was purchased from Chemical Procurement The Journal of Physical Chemistry
Laboratories, Inc. and was degassed and distilled under high vacuum before use. The anion radicals were pre(1) Part I : I?. J. Smentowski and G. R. Stevenson, J . Amer. Chem. Roc., 89, 5120 (1967). The present article is d s o part I1 of the series Anion Radicals in Liquid Ammonia (Part I : I?, ,J. Smentowski and 0.R. Stevenson, ibid., 90, 4661 (1968)). (2) (a) E. Grunwald, Anal. Chem., 26, 1696 (1954); (b) S.Winstein, E. Clippinger, A. H. Fainberg, and G. C. Robinson, J . Amer. Chem. Soc., 7 6 , 2697 (1954); (c) S. Winstein and 0. C. Robinson, ibid., 80, 169 (1958). (3) (a) N. M. Atherton and 5. I. Weissman, J . Amer. Chem. SOC.,83, 1330 (1961); (b) P. J. Zandstra and S. I. Weissman, ibid., 84, 4408 (1962); (c) F. C. Adam and 8. I. Weissman, ibid., 80, 1518 (1958). (4) (a) A. H. Reddoch, J . Chem. Phys., 43, 225 (1965); (b) N. Hirota and R. Kreilick, J. Amer. Chem. SOC.,88, 614 (1966); (c) R. Chang and C. 8. Johnson, ibid., 88, 2238 (1966); (d) N. Hirota, ibid., 90,3603 (1968); (e) N. Hirota, R. Carraway, and W. Schook, ibid., 90, 3611 (1968); A. M. Hermann, A. Rembaum, and W.R. Carper, J . Phys. Chem., 71, 2661 (1967). (5) (a) P. Chang, R. V. Slates, and M. Sewarc, J . Phys. Chem., 70, 3180 (1966); D. N. Bhattacharyya, C. L. Lee, J. Smid, and XI, Sswarc, ibid., 69, 112 (1965); (c) C. Carvajal, K. J. Tolle, J. Smid, and M. Sswarc, J . Amer. Chem. SOC.,87, 6548 (1965). (6) (a) T. E. Hogen-Esch and J. Smid, ibid., 88, 307, 318 (1966); (b) T. E, Hogen-Esch and J. Smid, ibid., 87, 669 (1965); ( e ) J. Smid, ibid., 87, 665 (1965). (7) (a) N. Hirota, J . Phys. Chem., 71, 127 (1967); (b) M. C. R. Symons, ibid., 71, 172 (1967); ( 0 ) N. Hirota, J . Amer. Chem. Soc., 89, 32 (1967).
(8) F. J. Smentowski and G . R. Stevenson, J . Amer. Chem. Soc., 89, 6120 (1967). (9) H. L. Strauss, T. J. Kats, and G. K. Fraenkel, ibid., 85, 2360 (1963). (10) (a) F. J. Smentowski and G. R. Stevenson, ibid., 90, 4661 (1968); (b) H. J. Chen and M. Bersohn, Mol. Phy6. 13, 573 (1967). (11) (a) D. H. Levy, Ph.D. Thesis, University of California. 1965; (b) D. H. Levy and R. Myers, J. Chem. Phys. 41, 1062 (1964); 42, 3731 (1965); 43, 3063 (1965); 44, 4177 (1966).
TEMPERATURE-DEPENDENT ELECTRON SPIN RESONANCE STUDIES
341
pared by reduction of COT with the alkali metal-liquid ammonia solution.*0a Activation energies were determined from the slope of In (line width) vs. 1/RT plots.12 Line widths were measured between the extrema in the first-derivative spectrum. In the linear region of the In (line width) vs. 1/RT plot, all hyperfine components of the same system gave the same line width. Spectra of each system were run at several different metal and COT concentrations to verify that the energy of activation for the line broadening is independent of the concentraFigure 1. Plot of In (line width) us. 103/RT for the system tion of the dianion. The concentration of the dianion COT-NHs-Li (0.05 M COT-0.05 M Li) in the region was low enough that a visible amount of the dianion of fast exchange. salt did not precipitate out of solution. For the system COT-NHa-K, kinetics for the line-broadening process were determined as described previously.* Equilibrium constants for the disproportionation (reaction 3) were measured by comparison of the spin concentration of the ammonia system to that of the system COTTHF-Li, where the equilibrium constant is k n ~ w n . ~ , ’ ~ The spectra were recorded using the X band of a Varian V-4502-15 esr spectrometer with a 12-in. magnet. Temperature was controlled within & l oby a Varian V-4557 variable-temperature controller. A copper-constantan thermocouple was used to caliFigure 2. Plot of In (line width) us. 103/RT for the system brate the variable-temperature controller. Coupling COT-NH8-Na (0.136 M COT-0.060 M Na) in t,he region of fast exchange. constants and line widths were taken directly from the calibrated chart paper. 1i l
Results Three systems (COT-NH3-me.Id) were studied, where the metal is lithium, sodium, or potassium. Ion pairing affects the esr spectra of I, even in liquid ammonia, influencing the line widt,hs of the individual hyperfine splittings, the activation energy of the linebroadening process, and the spin concentration. For all systems studied, successive additions of metal to M COT leads to a wider line in the slow-exchange region and a narrower line in the fast-exchange region. Because of precipitate formation, spectral measurements were not possible for > M concentrations of J and 11. Over the linear range of the In (line width) vs. 1/RT plot (Figures 1-3) the same activation energy is obtained for different amounts of metal added. For the three systems it was found that the activation energy of the line-broadening process over the linear range of the In (line width) vs. 1/RT plot was independent of the dianion I1 concentration, The dependence of the hyperfine line width of the cyclooctatetraenide salts is Li+ > Na+ > E(+ when the spin concentration in these systems is i 10-6 M . Figure 4 is a plot of the relative spin concentration vs. T,varying the counterion. Table I gives the equilibrium constants of the disproportionation (reaction 3) for the three systems a t 22”. For the COT-”8-I< system, the radical(s) ob-
I
/
04P 1
18
_LI _
-
-1.
1 ---A
22
20
JQ3 RT
-
,
24
Figure 3. Plot of In (line width) vs. lO3/Rl’ for the system COT-NHa-K (0.055 M COT-0.06 M K ) in the region of slow exchange.
served are dependent upon the ratio of COT t o K. No metal splitting is observed in any of these systems. COT-NH&i. Fast-exchange conditions (one line) prevail a t room temperature for all lithium systems, except for very low concentrations (approximately M ) of lithium when nine lines can be observed. The activation energy for those factors affecting the line width in this system is 1.1 kcal/mol over the temperature range -70 to 10”. The spin concentration of this system increases by a factor of 2.2 over the inter(12) (a) R. L. Ward and s. I . Weissman, J . Amer. Chem. SOC., 79, 2086 (1957); (b) M. T. Jones and S. I. Weissman, ibid., 84, 4269 (1962); (c) W. L. Reynolds, J . Phys. Chem., 6 7 , 2866 (19153); (dj E.deBoer and C. MacLean, J. Chem. PhUs., 44, 1334 (19GG). (13) T. J. Kats, J . Amer. Chem. SOC.,82, 3784,3785 (19130). Volume 73, Number 4
Februniu 1960
F. J. SMENTOWSKI AND GERALD R. STEVENSON
342
served for the lithium or sodium systems. K, for the disproportionation (reaction 3) for the potassium system is (1.4 1) X Table I1 gives a representative series of systems used to calculate K,. Two different radicals are observed for this system for the concentration of I1 from lo-* to 5 X M . At, the lower limit of dianion concentrations, I was not observed by esr spectroscopy. At the upper limit, I1 precipitated from solution. For high concentrations of metal ((K) > 6(COT)) radical 1 (UK = 3.28 0.01 G) is found. For low concentrations of metal ((K) < '/2(COT)), radical 2 ( a = ~ 3.30 f 0.01 G) is found. Table 111summarizes this information on the two radicals. Between these limits, the two radicals exist, superimposed on each other (Figure 5 ) . In both radicals 1 and 2 all nine lines are equivalent.
*
*
-70
-50
-30
-10 +IO T 'C----+
+30
+50
Figure 4. Plot of temperature us. the relative intensity of the esr signal for liquid NHa systems: 0, K reduction; X N a reduction; A, Li reduction. Intensity is the esr line height times the line width squared.
Table I : Equilibrium Constants a t 22' for the Disproportionation for Various COT Alkali Metal Systems Solvent
Metal
NHs
K
"a
Na Li
8"
THF THF DMF DMSO a
K Li Electrolytic" Electrolytic"
K89
Ref
( 1 . 4 i 1) X 10-8 (2.5 i: 2) X 10-2 (7 & 5) x 10-1 3 . 4 x 10-9 9 x 10-9 104
This work This work This work 9 9 18 18
lo4
Table I1 : Disproportionation Constant for the COT-NHs-K System, 22' (Representative Data) Sample
[COTla
10-6[COT* -1"
[COTa-la
I I1 111
0.04 0.10 0.038
4.9 4.2 1.9
0.05 0.038 0.004
10*Ke
1.2 0.46 2.5
Molar equilibrium concentration.
Carrier electrolyte, te;tra-n-propylammoiiium perchlorate.
val - 70 to 30". I(, for the disproportioiiatioiI (reaction 3) for the lithium system is (7 f 5 ) X lO-I, greater than the sodium or potassiimi system^. COT-NHB-Nu. There are several activation energies observed for those factors sff ecting the line width in this system (Figure 2). The activation energy for the lower temperatures is 0.50 kcnl/mol. At the higher temperatures, a "negative" activation energy of 1.0 kcal/mol is observed. Fast-exchange conditions prcvail for sodium systems when the dianion concentration is greater than 111. At lower concentrations of the dianion 11, nine lines can be observed. The spin concentration of this system increases by a factor of 2.G over the interval -70 to 30". K , for thedisproportionation (reaction 3) for the sodium system is (2.5 f 2) x 10-2. COT-"8-K. For this system, nine lines are always observed, even for high (e.g., 0.05 M ) concentrations of dianion 11. The activation energy for those factors affecting the line width in this system is 0.60 kcal/mol. Comparing the line widths of the potassium system with the other metal systems, the line for the potassium salt is much narrower than the line for the sodium or lithium systems, The spin concentration of this system increases by a factor of 8 over the interval 70 to 30". This variation is larger than that ob-
-
The Journal of Physical Chemistry
Table 111: Dependence of the Radical(s) Observed on the Ratio of COT to Its Dianion for the COT-NHs-K System a t 22" [COTla/ [KIb
[COTI"
fK'lc
[+-I"
fnlC
1.3 7.7 0.9 1.1 8.3 1.4 2.2 4.4 0.6 8.8 8.0 16.3 23.7 1.1 4.9
0.30 0.10 0.053 0.033 0.048 0.083 0.103 0.203 0.303 0.403 0.064 0.13 0.19 0.013 0.19
0.20 0.013
0.10
0.061 0.030 0.058
0.03 0.015 0.020 0.030 0.023 0.023 0,023 0,023 0.004 0.004 0.004 0.006 0.020
0.20 0.099 0.023
0.060 0.046 0.046 0.046 0.046 0 * 008
0.008 0.008 0.012 0 039 I
0.0065
Radical absd
1 2 f
0.018
1
0.019 0.053
2 1
0.080
1, 2
0.180 0.280 0.380 0.060 0.13 0.19 0,012 0.17
1, 2 2 2 2 2 2 1 ' 1
" Initial COT concentration, M . Initial K concentration, g-atoms/l. Equilibrium concentration, M .
For the system COT-NHsKJ kinetics for the linebroadening process were attempted. However, the concentration of I1 could not be varied over a large enough range to obtain meaningful results for either radical 1 or 2.
343
I
^
1
. . /
!
j0.421I , !
3a a G 2
Na+ > K+. In the ether solvents DME and T H F the reverse order prevails*: K + > Na+ > Li+. Conductance studies1° of anion radicals indicated this same relative order of cation migration in the ether solvents, with the small cation coordinating to the solvent to form an immobile ion. Also for systems undergoing the electron transfer (eq 5 ) , it appears this transfer is faster for the dissociated ion than for the ion-paired anion radicaLZ0 However, addition of water to aprotic solvents has in some instancesz1 led to a decrease in the electron-tr ansfer rate. At this point, further details of the electron transfer for these cyclooctatetraene systems cannot be given. Electron transfer for simpler systemszzis complex. What is most surprising are the results obtained for the equilibrium constants for the disproportionation (eq 3), given in Table I. All three systems, lithium, sodium, and potassium, give higher concentrations of anion radical I than the ether solvents T H F and DME. Comparing the two systems COT-THF--Li and COTNH3-Li, the Ke‘s differ by 100,000,000! The anion radical I appears to be favored over the dianion I1 for the system COT-NH3--electr~de,11but the equilibrium constant was not determined. Allendoerfer and Riegerl*have observed a similar dramatic shift in K , of the disproportionation (eq 3) for the dissociated ions for the systems COT-DMF-electrode and COT-DMSO-electrode. They suggest that the shift in equilibrium (eq 3) in going from Dl\IF and DAIS0 to ethereal solvents is due to the formation of a strong alkali metal-dianion complex. For the six systems COTsolvent--J\r, where solvent = THF, DATE, or NHa and ib1 = Li or Na, an increase in the dielectric constant of the solvent (i.e,, lowering of the temperature of a particular system or changing the solvents) together with a small cation favors the anion radical. Similarly, from Table I, the concentration of I is greater for those syst,ems with solvents of high dielectric constant. In solvents of low dieleclric constant, aromatic hydrocarbon dianions are more highly associated with alkaIi (16) I?. J. Smentowski and G. R.’Stevenson, submitted for public& tion in J . Amer. Chem. SOC. (17) (a) W. L. Jolly and C. J. Hallada, “Non-aqueous Solvent Systems,” T. C. Waddington, Ed., Academic Press, New York, N. Y., 1965, p 1; (b) G. Lepoutre and M. J. Sienko, “Metal-Ammonia Solutions,” Benjamin, New York, N. Y., 1964. (18) R. D. Allendoerfer and P. H. Rieger, J. Amer. Chem. Soc., 87, 2336 (1965). (19) (a) R. V. Slates and M. Szwarc, J . Phys. Chent., 69,4124 (1966); (b) P. Chang, R. V. Slates and M. Szwarc, ibid., 70, 3180 (1966). (20) (a) R. Chang and C. S. Johnson, Jr., J. Amer. Chem. SOC.,88, 2338 (1966); (b) J. E. Harriman and A. H. Maki, J . Chem. Phys., 39, 778 (1963). (21) T. Layoff, T. Miller, R. N. Adams, A. Horsfield, and W. Procter, Nature, 205, 382 (1965). (22) (a) R. A. Marcus, “Annual Review of Physical Chemistry,” Vol. 15, H. Eyring, C. J. Christensen, and H. 8. Johnston, Ed., Annual Reviews Inc., Palo Alto, Calif., 1964; (b) R. A. LMarcus, J. Chem. Phys., 43, 2654 (1965).
TEMPERATURE-DEPENDENT ELECTRON SPINRESONANCE STUDIES metal ions than the corresponding monoanions.28 No further relationship between the above-mentioned factors and the thermodynamic function^^^^^ affecting K , can be made. The greater stability of the COT dianion I1 over the anion radical for systems such as COT-DME-Li has been widely quoted as an application of the Hiickel4n 2 rule. However, the system COT-1\;H3-Li gives approximately equal concentrations of anion radical I and the dianion 11. HMO calculations show the anion radical I and the dianion I1 have the same resonance energies. Approximate SCF-MO calculations25 using the Pople-Santry-Segal methodz6 show the anion radical I to be much more stable, 110 kcal/mol, than the dianion 11. However, these calculations do not consider solvation effects.
+
Summary and Conclusions Even in ammonia, a solvent having a reasonably high dielectric constant, ion pairing is a major factor in the variation of some of the physical properties of anion radical I. For the alkali metal systems in liquid ammonia, the major source of line broadening appears to be electron transfer from dianion I1 to anion radical I. The activation energy of the electron transfer between dianion I1 and anion radical I is not as dependent upon the counterion as it is in either solvents. The line widths of the individual hyperfine splittings are a logarithmic function of the temperature (Figures 1 and 2) and are somewhat dependent upon the counterion. Again, the variation of spin concen-
345
tration with temperature, solvent, and counterion is indicative of the influence of ion pairing on the disproportionation (eq 3). The dielectric constant of the solvent and its coordinating ability for the cation are important in determining the rate of electron transfer between the anion radical I and its dianion 11. These factors are also important in determining the direction of the disproportionation (eq 3).
Acknowledgment. The authors are indebted to Professor B. J. Zwolinski for helpful discussion. The authors thank the National Aeronautics and Space Agency for Fellowship support to G. R. Stevenson. The esr spectrometer was made available by the National Science Foundation Grant GP-3767. Support from the Research Council of Texas A & TVI University is gratefully acknowledged. Time for machine calculations was made available by the Data Processing Center of Texas A & M University. Acknowledgment is made to the donors of The Petroleum Research Fund, administered by the American Chemical Society, for partial support of this research. (23) (a) G. J. Hoijtink and P. H. van der Meij, 2. Phys. Chem. (Frankfurt am Main), 20, 1 (1959); (b) G. J. Hoijtink, E. de Boer, P. H. van der Meij, and W. P. Weijland, Rec. Trav. Chim., 75, 487 (1956). (24) G. R. Stevenson, Ph.D. Thesis, Texas A & M University, 1969. (25) F. J. Smentowski, B. D. Faubion, and R. M. Hedges, submitted for publication in J. Chem. Phys. (26) (a) J. A. Pople, D. P. Santry, and G. A. Segal, J. Chem. Phys., 43, S129 (1965); (b) J. A. Pople and G. A. Segal, ibid., 43, 8136 (1965); (c) J. A. Pople and G. A. Segal, ibid., 44, 3289 (1966).
Volume ?3?Number 8 February 1069
