E. G. HOHN,J. A. OLANDER, AND M. C. DAY
3880
COScould be produced by the photolysis (9) of OSin a COzmatrix. They proposed that (9)
COSwas produced by the reaction of atomic oxygen from with COz. In the present study it was suggested that a limiting factor to O F radical production may be photolysis of the radicals as shown in eq 6 to give oxygen and fluorine atoms. If the generated oxygen atoms’ react (eq 10) with C02 to give COS then OF radicals could be regenerated by the reaction (eq 11) 0 3
+ coz +co, F + COS+OF + COz 0
(10) (11)
of atomic F with COS. In this manner, CO:, would be acting as an oxygen transfer agent and would offset the photolytic loss of OF radical by allowing for the reformation of OF as shown in eq 10 and ll. If this proposed mechanism is valid, then the magnitude of the observed improvement (run 3, Figure 1) in the presence of COZ suggests that photolytic loss of OF (eq 6) contributes significantly to total OF losses. The absence of observed infrared absorptionse for COSin the COZrun was not considered as negative evidence for their presence in view of their transient existence in this system. An alternate function for COSmust also be considered because of the previous observatione that COa itself is subject to photolytic decomposition to give 0 atoms and
Ion-Solvent Interactions.
C02. The overall effect of improved OF radical production would be the same as in the previous mechanism; however, in this case it would be unnecessary to invoke the F COSreaction. The improved OF radical production would occur because the reversible reaction
+
hu
O
+ COa C,COS
(12) would increase the range of migration of 0 atoms within the matrix. Thus, in the absence of C02the generated 0 and F atoms from O F decomposition would migrate to isolated matrix sites. However, in the presence of GO, throughout the matrix, reaction 12 could increase the range of 0 atom migration and thereby increase the probability of an 0 F reaction to regenerate OF. OF Radical Bond Energy. The observation of OF radical formation during photolysis of F2-Nz0 mixtures confirms the presence of this intermediate in the forrnation of OFZ,as previously proposed by Ogden and Turnern4 It also supports their lowerlimit estimate of about 40 kcal/mol for the OF radical bond energy based on the energy requirements for reaction 2. Acknowledgments. The author wishes to thank Drs. P. H. Lewis and S. A. Francis for helpful discussions and Mr. G. H. Post for his help with the experimental portion of the work.
+
(7) As suggested by a reviewer it should be noted that eq 10 assumes the dissipation of hot 0 atom energy in the formation of COa.
Infrared Studies of Solvation of the Sodium Ion1i2
by E. G. Hohn,8J. A. Olander, and M. C. Day Department of Chemistry, Louisiana State University, Baton Rouge, Louisiana 70805 (Received April 14, 1969)
The specific solvation of the sodium ion by tetrahydrofuran has been observed by infrared techniques. Unperturbed voo0bands occur at 1071 and 913 cm-1. In the presence of the sodium ion, new bands occur at -1053 5 respectively, indicating complexation of the sodium ion. The band intensities are observed to and ~ 8 9 cm-l, be dependent on the ratio of ether: salt. This is used to approximate stability constants for the stepwise complexation of the sodium ion by tetrahydrofuran.
Introduction The importance of ion-solvent interactions on ionic conductance and extent of ion pairing was clearly pointed out by Gilkerson4 and has been extensively studied by several research groups.5-9 More recently there has been considerable interest in the effects of solvent interactions on the rates of anionic polymerization reaction^,^^-^^ but in Spite Of the large amount Of reThe Journal of Physical Chemistry
search in this area, there are yet numerous unanswered questions concerning the nature of the solvated species (I)Reprint requests should be sent to M. C. Day. (2) Presented in part at Southwestern Meeting of the American Chemical Society, Dec 1967. (3) Department of Chemistry and Chemical Engineering, University of Saskatchewan. Saskatoon, sask, (4) R. Gilkerson, J . Chern.Phys., 25,1199 (1956).
w.
ION-SOLVENT INTERACTIONS and the effect solvation may have on the type of ion pair present in solution. The stability of a complex between a cation and a Lewis base should be dependent on the size of the cation and the basicity of the coordinating species. I n order to obtain a strong ion-solvent interaction, it is necessary to have a salt with a small cation. It is also desirable to be able to control the concentration of the coordinating species. This requires the use of a solvent that is inert, or at least, a relatively poor coordinating agent to which the coordinating species can be added in controlled amounts. Unfortunately, salts containing small cations are not generally soluble in noncoordinating solvents. As a consequence, in the use of mixed solvent systems, the study of the complexation of a small cation by a coordinating species is normally complicated by competition with the solvent. Recently, we have shown that the specific solvation of the cation can be studied in mixed systems by nmr14 and conductancel5 methods. Sodium tetrabutylaluminate (NaA1Bu4) has the unique property of being soluble in saturated hydrocarbon solvents, and it was shown that there is no significant interaction between the NaAlBu4 and the hydrocarbon s01vent.l~ The solvent can thus be considered to act essentially as a dispersing medium for the salt. This overcomes the problem of competitive solvation normally found in mixed solvent systems, and the effect of the controlled addition of a complexing species can then be followed. We wish to report here the use of infrared spectroscopy in the study of the specific complexation of the sodium ion by tetrahydrofuran (THF).
Experimental Section Infrared spectra were recorded using a Beckman IR-7 spectrophotometer and Beckman R!todel FH-01 vacuum-tight liquid cells equipped with potassium bromide windows. Amalgamated lead spacers were used to give a path length of 0.03 mm. The spectra were recorded a t ambient temperature of ca. 25”. Preparation of Sodium Tetra~utyZuZum~nate.NaA1Bu4 was prepared by direct addition of 1 mol of aluminum tributyl (198 g) t o a sodium dispersion containing an excess of sodium (23 g) in 200 ml of n-heptane. The mixture is stirred and refluxed under a nitrogen atmosphere for 2 hr and then filtered through a finefritted glass funnel in a nitrogen atmosphere drybox.15 The n-heptane is removed from the NaAlBu4 by vacuum evaporation, after which the product is recrystallized from n-pentane at Dry Ice temperatures. The white, crystalline NaA1Bu4 is then dried under vacuum. Sodium tetraethylaluminate (NaA1Et4) and sodium tetraoctylaluminate (NaAloctyld) were prepared by the general method described for NaAlBu4 with toluene used as a reaction solvent rather than heptane and the
3881 appropriate aluminum alkyl, Alocty13 or A1Et3, used in place df AlBu3. Solvents. Cyclohexane was treated with an HzSO4HK03 mixture, dried over KOH, and distilled from CaH2. THF and BuzO were dried over KOH and distilled from LiAlH4. Distillation and collection of solvents were carried out under a nitrogen atmosphere. Preparation of Solutions. Solutions of varying T H F concentration and a constant NaA1Bu4 concentration were prepared by adding an aliquot of a standard NaA1Bu4 solution in cyclohexane to a weighed portion of T H F and diluting to volume with cyclohexane. All manipulations involving exposure of NaAlBu4, solvents, or solutions were carried out in a nitrogen drybox. Experimental Procedure. I n order to determine the concentration of unperturbed T H F in NaA1Bu4-THFcyclohexane solutions, spectra of solutions of T H F in cyclohexane of a known THF concentration were recorded from 900 to 1150 em-’. The intensity of the 1071-cm-1 band of T H F in these solutions was obtained by constructing an artificial base lineI6 and using it to obtain a value for the incident radiation (%To). The per cent transmittance (%T) was read from the point of maximum absorption for the band. These per cent transmittances were then converted to the respective absorbances (A0 and A ) , and their difference ( A - Ao) was plotted os. concentration of THF. Linear plots were obtained in all cases. The concentration of unperturbed T H F in NaAlBudTHF-cyclohexane solutions was then found by determining an ( A Ao)value for the 1071-em-l band as was done for the standard solutions. This value was then used to determine the concentration of unperturbed THF from the standard plot. Some error is inherent in this method since the -1053- and 1071-~m-~ bands are not completely separated; thus, the intensity of either band will be enhanced due t o overlap with the other. This error is considered to be a minor one and
-
(5) R. M.Fuoss, et al., J . Phys. Chem., 69, 2576 (1965),and many others in the series. (6) C. A. Kraus, J.Chem. Educ., 35,324 (1958). (7) W.R. Gilkerson and J. B. Ezell, J . Amer. Chem. SOL,89, 808 (1967). (8) T.E.Hogen-Esch and J. Smid, ibi;d., 88,318 (1966). (9) A. D’Aprano and R. Triolo, J . Phys. Chem., 71,3474 (1967). (10) T.Shimomura, J. Smid, and M. Szwarc, J . Amer. Chem. SOC., 89, 5743 (1967). (11) 3. E. L. Roovers and S. Bywater, Can. J . Chem., 46, 2711 (1968). (12) F.A. Dainton, et al., Makromol. Chem., 89, 257 (1965). (13) 3. E.L.Roovers and S. Bywater, Trans. Faradag SOC.,62, 701 (1966). (14) E. Schaschel and M. C. Day, J . Amer. Chem. SOC.,90, 503 (1968). (15)C.N.Hammonds and M. C. Day, J . Phys. Chem., 73, 1151 (1969) (16) H. H. Willard, L. L. Merritt, Jr., and J. A. Dean, “Instrumental Methods of Analysis,” 4th ed, D. Van Nostrand Co., Inc., Princeton, N. J., 1965.
Volume 7.9, Number 11
November 1969
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_-
3882
E. G. HOHN,J. A. OLANDER, AND M. C. DAY I
I
I
I
I
I
1
t
I
I
I
I
I
I
0 . 2 4 5 M NaAIBu,
90 T
0.269M N a AI Bu, R a t i o f T H F,l/N 0 995
;9"
O
1100
V i
a AI B u4]
----.
1000
0
900
CG' Figure 1. Infrared spectra of NaAlBurTHF-cyclohexane solutions of constant salt concentration and varying T H F concentration.
relatively unimportant within the experimental procedures used. To ensure that cell constants did not change in the course of a study, alternating series of spectra of standard THF-cyclohexane and NaA1Bu~-THF-cyclohexane solutions were obtained.
Results and Discussion The infrared spectra of NaAlBu4-THF-cyclohexane solutions of constant NaAlBu4 concentration and varying T H F concentration in the 900-1150-~m-~region are shown in Figure 1. The asymmetric C-0-C stretching vibration of T H F appears at 1071 cm-I while the symmetric C-0-C mode appears at 913 cm-1.18 Bands a t 902, 1015, and 1040 cm-' are due to cyclohexane. The AIBu4- anion shows broad absorptions centered a t 1060 and 1145 cm-', and moderately intense absorptions at 992 and 955 cm-'. As can be seen from Figure 1, the 1071- and 913-cm-' bands of T H F are not evident a t T H F : salt ratios below 1:1, but a new band appears at -1048 cm-'. A second new band appears in the 900-cm-' region but is masked by the 902-cm-1 band of cyclohexane. As the T H F : salt ratio is increased past the 1 : l ratio, the band a t -1048 cm-l shifts towa,rd a final value of 1053 ern-' and increases in intensity. Correspondingly, the 1071om-' band becomes evident, and the band at 913 cm-' appears as a shoulder on the 902-cm-' band of cyclohexane, thus indicating the presence of free THF. The growth of the band at -902 cm-l with an increase in T H F concentration can be attributed to the overlap of the perturbed C-0-C symmetric T H F band and the 9 0 2 - ~ m -cyclohexane ~ band. The complexation of THF to group,III halides has been shown to shift the frequencies of both the symmetric and asymmetric C-0-C stretching modes of The Journal of Physical Chemistry
I
1
2.0
I
4.0
I
I
6.0
1
I
8.0
I
1
I
10.0
1
i I
12.0
Figure 2. Average value plots, li, of the ratio of bound THF: salt us. ratio of total THF: salt for NaAlBur concentrations of 0.245 and 0.269 M .
T H F to lower v a 1 ~ e s . l ~The new band at -1053 cm-l is thus considered to result from the C-0-C stretch of T H F molecules complexed with the sodium ion. Spectra of NaAlEt4 and NaAlocty14 in T H F show essentially the same perturbed bands observed with the NaAlBu4-THF-cyclohexane system, thereby substantiating the interpretation that the complexation occurs with the sodium ion. Other similar complexing agents should exhibit analogous behavior. This was shown to be true for di-n-butyl ether which shows a shift from the unperturbed vcoc of 1125 cm-l to a perturbed value of 1083 cm-I upon complexation. As can be seen from Figure 1,the absence of the 1071 cm-' band at ratios of T H F : salt less than 1: 1, and the relative intensity changes as the T H F : salt ratio is varied indicate an equilibrium between the 1 : l complex and additional THF molecules. I n Figure 2, average value plots ( E ) of the ratio of (bound THF): salt us. (total THF) :salt ([THFt,] : [salt] us. [THFt]: [salt]) are shown for salt concentrations of 0.245 and 0.269 M . Below a ratio of 1:1, within experimental limitations, all of the T H F is observed to be in the complexed form. The 1: 1 complex is then in equilibrium with additional THF molecules, giving a limiting average value of 4, indicating a maximum coordination of 4 for T H F and sodium ion under these conditions. This can be considered as NaaTHFf, A1Bu4-
+ 3THF JcNa*4THF+,AlBu4-
with the exact role of the anion or nature of the species (17) G . M. Barrow and S. Searles, J. Amer. Chem. Soc., '75, 1175 (1953). (18) N.B. Colthup, L. H. Daly, and S. E. Wiberley, "Introduction t o Infrared and Raman Spectroscopy," Academic Press, New York, N.Y.,1964,p 272. (19) J. Lewis, J. R. Miller, R. L. Richards, and A. Thompson, J. Chem. SOC., 6850 (1965).
ION-SOLVENT INTERACTIONS
3883
Table I : Approximate Stepwise Stability Constants for (Na.zTHF)+ Salt ooncn, M
kl
ke
ka
kr
0.269
Large Large
28
4.6
33
5.5
0.9 0.9
0.245
(e.g., solvent-separated or contact ion-pair, etc.) not known. Using Bjerrum’s methodz0of plotting 5 vs. the negative logarithm of free ligand concentration (p [THFI]), approximations to the stepwise stability constants for the THE” complex with the sodium ion have been made. These are listed in Table I for both the 0.245 and 0.269 M solutions of NaAlBu4. These values are only first order approximations to the stability constants but do give representative orders of magnitude. The discrepancies between the values for the two concentrations are on the order of the accuracy of the system.
From the above, it can be concluded that the sodium ion forms a very stable 1 : l complex with THF, and this is in equilibrium with three additional T H F molecules forming a four-coordinated species. Previous studies of this system using nmr14 and conductance16 techniques have led to an analogous, but less definitive, interpretation. Although there may be a concentration dependence as we11 as solvent dependence of the solvation number of the sodium ion, it would seem that the maximum coordination number of 4 is now established for this particular system a t a salt concentration in the region of 0.25 M.
Acknowledgment. Support of this work by National Science Foundation Grants G P 1967 and GP 6421 and a National Aeronautics and Space Administration Traineeship for J. A. Olander is gratefully acknowledged. (20) J. Bjerrum, “Metal Ammine Formation in Aqueous Solution,” P. Haase and Son, Copenhagen, 1941.
Volume 78, Number 11
November 1969
