J . Phys. Chem. 1986, 90, 4684-4686
4684
GENERAL PHYSICAL CHEMISTRY Dielectric Properties of Solutions of Some Onium Salts In Dichloromethane B. Gestblom, Department of Physics, University of Uppsala, S- 75127 Uppsala, Sweden
I. Svorsterl, and J. Songstad*
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Department of Chemistry, University of Bergen, N-5000 Bergen, Norway (Received: August 6, 1985; In Final Form: March 18, 1986)
A study of the concentration dependence of the dielectric properties of solutions of R4NC104(R = Et, Pr, Bu, Hex, and Dec), Ph4AsI, and [(Ph3P),N]C1O4in dichloromethane is presented. The results are discussed in relation to findings from conductivity measurements on the same electrolyte solutions.
When an ionic compound, A+B-, is dissolved in a solvent of fairly low static permittivity, a number of different species may be present in the solution: free and/or solvated ions, A+ and B-;l contact ion pairs, [AB], and/or various types of solvent-separated ion pairs, [AIB];*triple ions, [ABA]’ and [BAB]-;3 quadrupoles, [AB],; and/or higher aggregates, [AB],.4 For concentrations that are of interest for most practical purposes, the individual concentrations seem to vary in a most complex way with the amount of dissolved salt, the size and structure of the ions, the structure and permittivity of the solvent, etc. Numerous conductivity studies have been performed in this class of solvents and it is well-known that the molar conductivity, A, above a certain concentration is far higher than that calculated from the association constant, KA, as determined in extremely dilute In dichloromethane this discrepancy between calculated and experimental molar conductivities starts to be M.’ This phenomenon was originally significant at 1 X explained by Fuoss and Kraus3 to be due to the formation of charged triple ions from nonconducting ion pairs. Cavell and Knight: however, proposed that the permittivity of the medium rather than that of the pure solvent is the determining factor and derived an expression for the association constant, KA,as a function of the permittivity of the solution based upon the original equation by Bjenum? The increasing static permittivity of the will lower KA, leading to a larger fraction of free or solvated ions.8 In this paper we report on a systematic study of the concentration dependence of the dielectric properties of solutions of some onium salts, mainly perchlorates, in dichloromethane at 20 OC.
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(1) Parker, A. J. Chem. Rev. 1969, 69, 1. (2) Winstein, S.;Kleindienst, P. E., Jr.; Robinson, G. C. J . Am. Chem. Soc. 1961, 83, 885. (3) Fuoss, R. M.;Kraus, C. A. J . A m . Chem. SOC.1933, 55, 2387. (4) Batson, F. M.; Kraus, C. A. J. Am. Chem. SOC.1934, 56, 2017. (5) Bekkevoll, S.; Svorstol, I.; Hailand, H.; Songstad, J. Acta Chem. Scand., Ser. B 1983, 37, 935. (6) Walden, P. Z . Phys. Chem., Stoechiom. Verwandtschaftsl. 1902, 30, 513; 1922, 100, 512; 1930, 147, 1. (7) Svorstd, I.; Songstad, J. Acta. Chem. Scand., Ser. B 1985, 39, 639. (8) Cavell, E. A. S.; Knight, P. C. 2. Phys. Chem. (Neue Folge) 1968,57, 331.
(9) Bjerrum, N. K . Dan. Vidensk. Selsk. Skr., Naturuidensk, Math. Afd. 1926, 7 (9).
(10) Lestrade, J.-C.; Badiali, J.-P.; Cachet, H. Dielectr. Relat. Mol. Processes 1975, 2. (1 1) Cachet, H.; Cyrot, A.; Fekir, M.; Lestrade, J.-C. J. Phys. Chem. 1979, 83,2419. Saar, D.; Brauner, J.; Farber, H.; Petrucci, S. J . Phys. Chem. 1980, 84, 341.
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Results and Discussion The permittivity spectra were determined in the range 0.05-10 GHz by the total transmission time domain spectroscopy (TDS) technique. The spectra showed two dispersions and were analyzed in terms of a two-relaxation-time model function
o is the conductivity, to is the permittivity of free space, and t, and E, are the limiting permittivities at low and high frequencies, respectively. The ts- t1 dispersion, with relaxation time T ~is, due to ion pairs in the solution, the other dispersion being due to the solvent itself. Further experimental details will be published.12 The results are shown in Figure 1, where the static permittivity, ts, and the relaxation time, T ~ are , plotted vs. the molar concentration of the dissolved salts. The permittivity of all solutions is seen to increase with concentration, which is in accordance with the results from similar studies in solvents of low dielectric constant.lO*llAt the same time the t l value decreases with concentration due to the decreasing volume fraction of the solvent: from t 1 = 9.1 at infinite dilution to el = 8.5 for 0.316 M solutions of Bu4NC104and Hex4NC104,to c1 = 8.1 for a 0.219 M solution of Dec4NC104, and to t l = 8.8 for a 0.156 M solution of Ph4AsI. For solutions of Et4NC104up to 0.060 M and of Pr4NC104up to 0.167 M this decrease in t l was found to be negligible. For dilute solutions of R4NC104the permittivity increases approximately linearly with the concentration with slopes, dt,/dc, fairly independent of the size of the cation, from -40 M-’ for Et4NC104 to -50 M-I for Dec4NC104. When the ion pairs are assumed to be the only species of sufficient dipole moment to alter the permittivity of the solutions, this linearity in the e,-c plots indicates that the concentration of these species is proportional to the total concentration in dilute solutions. Furthermore, the similar slopes suggest that the dipole moments of the ion pairs cannot be greatly dependent on the size of the cation. This latter conclusion is in accord with the result from a recent conductivity study of several tetraalkylammonium perchlorates in dichloromethane, in which a Bjerrum treatment9 of the conductivity data showed that the ion pairs were of the contact type with distances of closest approach ranging from 4.9 (3) 8, for Et4NC104 to 5.8 (3) 8, for Dec4NC10k7 A viscosity study on %N+ salts in dichloromethane has similarly shown that R4N+cations with large R groups “curl
(12) Gestblom, B.; Songstad, J., to be published.
0 1986 American Chemical Society
The Journal of Physical Chemistry, Vol. 90, No. 19, 1986 4685
Dielectric Properties of Onium Salt Solutions 0.5
0.4
t-
14
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14.0
1
w"
12.0
11.0
10.0
9.0
0.1
0.2
0.3
0.4
0.5
Figure 1. The static permittivity, e,, and the relaxation time, T , , as a function of the concentration of onium salts in dichloromethane at 20.0 O C . (Circled points on curves for Pr4NC104refer to studies on Pr4NI.)
up" to "hard spheres" in this ~ o l v e n t . ' ~The longer relaxation times, T , , for the salts of the larger cations at low concentrations seem to indicate that these "hard spheres" loosen up to some extent when sufficient space is available. This loosening up of the "hard spheres" may only be considered as some type of conformational reorganization of the long alkyl groups. The ion pairs may still be considered as contact ion pairs in dilute solutions as concluded from the conductivity data.7 For higher concentrations the effect of the size of the cations is most pronounced. For the salts of the larger cations the e,-c plots level off rapidly. Interestingly, these are the salts for which only an inflection in the A-c'I2 plots was observed.' For the salts of the smaller cations the static permittivity of the solutions continues to increase but with steadily decreasing slopes. These are the salts for which a minimum and a maximum in the A-c'I2 plots were d e t e ~ t e d .These ~ observations seem to indicate that the permittivity of the solutions is the governing factor with regard to the conductivity of concentrated electrolytes in solvents of low e,, as was indeed proposed by Cave11 and Knight.8 Ph4AsI and [PNP]C104 seem to behave differently. For low concentrations the permittivity of the solutions increases strongly with concentration with a dt,/dc of -150 M-'. This indicates that the dipole moments of the ion pairs from these two salts are considerably larger than viewed by the ionic radii. In the conductivity study it was concluded that in very dilute solutions these two salts gave solvent-separated ion pairs since the calculated distances of closest approach were from 3 to 5 A larger than the sum of the van der Waals' radii.' The exceptionally long relaxation times for dilute solutions of Ph4AsI and [PNP]C104are in accordance with the results from the conductivity study. For conSvorstel, I.; Songstad, J. Acra Chem. Scand., Ser. E . , in press. (14) Svorstel, I.; Hoiland, H.; Songstad, J. Acta Chem. Scand., Ser. E (13)
1984, 38, 885.
centrations of Ph4AsI and [PNP]C104larger than -0.06 M the relaxation times are as for R4N+ salts of comparable size. This suggests a gradual transformation from solvent-separated ion pairs in dilute solutions to contact ion pairs in more concentrated solutions. The leveling off in the E,-c plots suggests that a maximum concentration of ion pairs is attained, a maximum that is strongly dependent upon the cation. This leveling off in the E,-c plots cannot be due to some redissociation to ions or to other conducting species; the molar conductivity of electrolytes in dichloromethane is known to decrease fairly rapidly for concentrations above 0.1-0.3 M, depending upon the size of the cation.' Furthermore, the near constancy at higher concentrations of the relaxation time points against the formation of new types of ion pairs. This suggests that the ion pairs are gradually consumed as quadrupoles or as higher aggregates. The E,-c curves may therefore be rationalized by the presence of two equilibria A+
+ B- e[AB]
2[AB]
KA
__ KM
[AB],
(2) (3)
in which KA is the usual association constant while KAArepresents the association constant for further association to quadrupoles or to other species of negligible dipole moment. Figure 1 shows that the maximum concentration of the ion pairs is lower for larger cations and will be attained at a lower concentration for salts of large cations than for salts of small ones. Apparently, while KA for R4NC104in dichloromethane is known to decrease slightly with increasing size of the cation7 but to be strongly dependent on the permittivity of the solution,8 KAA will increase strongly with increasing size of the cation. The net result will be that for concentrations well below the concentration at which the ion pair saturation concentration is attained, the salts
J . Phys. Chem. 1986, 90, 4686-4690
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of the larger cations will appear as the most dissociated salts. For higher concentrations, however, the salts of the smaller cations will provide a larger fraction of the ionic species. This conclusion is in agreement with the results from the conductivity study of R4N+ perchlorates in dichloromethane in the lO-’-IO-’ M concentration range.7 Since the A--C’/~curves for the various R4N+ perchlorates seem to intersect near the concentration by which a minimum or an inflection in these curves is observed, in diM,7 this critical concentration in chloromethane at -3 X solvents of low permittivity can probably be estimated with some certainty by means of Walden’s empirical relationship, e ~ , ~ ; ’ / ~ = con~t,63~ in which e represents the permittivity of the pure solvent. This reversal in the degree of association with the size of the cation of electrolytes in weakly dissociating media may prove to be of importance when the effect of cations upon rate constants in this class of solvents is to be considered; cf. a recent study by Eberson.’* It is generally assumed that triple ions in a solvent of low static permittivity will have much lower concentrations than other species like free or solvated ions, ion pairs, etc. The results obtained in the present study may thus not invalidate the concept of triple -~~~ ~~
~
~
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(15) Eberson, L. J . Am. Chem. SOC.1983. 105, 3192
ions. However, the fact remains that the conductivity of 0.1 M solutions of R,NC104 in dichloromethane is nearly 10 times the estimated values when assuming the ion pair association constants to be independent of the c~ncentration.~ Indeed, it appears to be most difficult to explain the e,-c and T ] - C plots if triple ions alone were responsible for the discrepancy. The various conductivity data in di~hloromethane,’*~~ together with the present results seem to suggest that an improved description of electrolytes in solvents of low permittivity can be based upon a series of equilibria between ions, ion pairs, and higher aggregates, as depicted by eq 2 and 3, rather than upon the involvement of triple ions. Given reliable expressions for the concentration dependence of KA and the cation dependence of KAA, one should be able to calculate with some confidence the concentrations of the various species in solutions of ionic compounds in dichloromethane. However, an extension to other solvents of low static permittivity does not seem possible unless the necessary conductivity and permittivity studies have been performed. Registry No. Et4NCI04, 2567-83-1; Pr4NC104, 15780-02-6; Bu4NC104,1923-70-2;Hex,NC1O4, 4656-81-9;Dec4NC104,62207-1 1-8; Ph4AsI,7422-32-4; [(Ph3P)2N]C104, 65300-04-1;dichloromethane, 7509-2.
Thermal Decomposition of Energetic Materlals. 16. Solld-Phase Structural Analysis and the Thermolysis of I,4-Dinltrofurazano[ 3,4-b ]piperazine Y. Oyumi, A. L. Rheingold, and T. B. Brill* Department of Chemistry, University of Delaware, Newark, Delaware 1971 6 (Received: March 13, 1986)
The solid-phase transition scheme above 293 K for the energetic molecule 1,4-dinitrofurazano[3,4-b]piperazine(DNFP) was determined by IR spectroscopy and differential thermal analysis to involve four polymorphs. The IR spectra and AH values from DTA suggest that differences in the puckering of the piperazine ring and degree of pyramidal distortion of the amine nitrogen atoms exist in some of the polymorphs while others are distinguished only by a very slight molecular redistribution in the crystal lattice. The crystal structure of one polymorph of DNFP is reported: monoclinic, P 2 , / c , a = 9.744 (2), b = 6.640 (2), c = 12.165 (2) A, p = 94.31 (2)O, Z = 4, RF = 4.40%, RwF= 5.20%. High-rate thermolysis of DNFP reveals that NO2 is the dominant decomposition product which is also predicted from u,(NOz) and the N-N bond distance correlations established in previous studies on nitramines containing the C2NN02unit. The thermolysis products of 1,4-dinitropiperazine and furazano[ 3,441piperazine are reported to help trace the origin of the products from DNFP.
Introduction Recent predictions1q2of high density and stored energy have fueled interest in nitraminofurazans. A particularly interesting, recently reported’ compound of this type is 1,4-dinitrofurazano[3,4-b]piperazine, 1, which is referred to as DNFP.
using differential thermal analysis, and the decomposition products in real time by using rapid-scan infrared spectroscopy. Aspects of the decomposition of DNFP are brought to light by comparing the products from DNFP to those of DNP, 2, and furazano[ 3,4,b]piperazine, 3, which are also characterized.
NO;.
NO2
NO2
No2
I
I
I 1
The efficacy of chemical propellants and energy-rich materials is closely linked to their solid-phase properties and their characteristics upon thermal decomposition. Therefore, we undertook a study of the molecular structure of D N F P by using X-ray crystallography and IR spectroscopy, the thermal properties by (1) Cichra, D. A.; Holden, J. R.; Dickinson, C. NSWC Report TR79-273, Naval Surface Weapons Center, Silver Spring, MD, 1980. (2) Rothstein, L. R.; Peterson, R. Propellants Explos. 1979, 4, 56. Rothstein, L. R. Propellants Explos. 1981, 6, 91. (3) Willer, R. L.; Moore, D. W. J . Org. Chem. 1985, 50, 5123.
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Experimental Section Purified samples of 1,2, and 3 were generously supplied by Dr. Rodney L. Willer (Morton-Thiokol, Elkton, MD).’ I R spectra of the DNFP polymorphs were obtained on samples produced by evaporation of solutions onto an NaCl plate or by heating and cooling the thin polycrystalline film of DNFP, as appropriate, in a homebuilt variable-temperature cell. A thermocouple in intimate contact with the sample and connected to a Fluke digital thermometer was used for temperature measurements. The temperature accuracy is estimated to be & l K. The cell was heated 0 1986 American Chemical Society
