NOTES
3366
appear to be unique to the isopropyl alcohol solvent system since the same order is observed in water and acetone as shown in Table 111.
Table I11 : pKi of Silver-Amine Complexes in Different Solvents
Solvent system
Div p K i of silver-amine complexelectric nIsoDi-n- Diiso- Tri-nconstant Butyl butyl butyl butyl butyl
Isopropyl alcohol, 18.3 9.17 8.75 8.13 6.84 5.70 0 . 5 M LiClOa Acetone, 0 . 5 M 20.7 10.29 9.68 10.05 8.83 . . . LiClOa Water 78.54 7.48 7.24 4.10 . . . 3.82
Acknowledgments. This work was financed in part by an Academic Year Extension of the Research Participation Program at Oklahoma State University sponsored by the National Science Foundation. The assistance and suggestions of Dr. Paul Arthur of Oklahoma State University are gratefully acknowledged.
Kinetics of the Gas Phase Pyrolysis
of Tetranitromethane' by J. M. Sullivan and A. E. Axworthy Rocketdyne, A Division of North Ameriean Aviation, Canoga Park, California (Received May 10, 1966)
Jonassen4 and others have suggested that the pKi of amine complexes should increase as the dielectric constant of the solvent decreases and in the case of ethylamine in the solvents water, ethanol, and isopropyl alcohol such a correlation appears to exist, perhaps because of the very similar nature of the solvents. That this is not the only factor involved can be seen by comparing the values of pKi in isopropyl alcohol, acetone, and water as shown in Table 111. Although it is often difficult to compare results of different investigations, the comparison shown in Table I11 would suggest that specific solvent effects must play a significant factor in complex stabilities although dielectric constant effects almost certainly are of importance. Further studies involving different solvent systems may be useful in determining the relative importance of these factors.
Conclusions Complexes of silver with amines fall into three groups, those with primary, secondary, and tertiary amines. Within these groups, the stability of the complexes correlates with the base strength of the amine. It does not, however, correspond to the order of the base strength of the amines in going from primary to secondary to tertiary, suggesting that steric hindrance is a significant factor. The relative stabilities of the amine complexes do not appear to change with changes in solvent although the solvent does affect the absolute values of instability constants. The dielectric constant may have a significant bearing on solvent changes on stability of complexes but specific solvent effects are also significant and may, in the case of solvents of quite different nature, be more important than the effects of dielectric constant. The Journal of Physical Chemistry
Reported herein are the results of our study of the pyrolysis of tetranitromethane (TNM).
Experimental Section The 90.6-ml monel electrically heated, stirred flow reactor was identical with that described by Sullivan and TNR4 vapor of 99% purity was mixed with helium and passed through the reactor at 1 atm total pressure and a TNM initial partial pressure of 1.09 mm (runs 1-23). The storage tank was repressurized with helium prior to run 24 giving an initial TIVM partial pressure of 0.48 mm in the final eight experiments. The reactor was by-passed to permit analysis of the unreacted TNM. A sample of the gas stream from the reactor (or the by-pass) was periodically analyzed for TNM by introducing it into the helium carrier gas stream of a gas chromatograph by means of an unlubricated Beckman gas sampling valve. The reduction in TNRI concentration with residence time in the reactor was followed chromatographically using a 3-ft column of 5% S.E.30 silicone oil on Celite, 60-100 mesh, at room temperature and a thermal conductivity detector. The flow rates through the reactor were measured with a soap bubble flow meter connected to the exit stream. The measured flow rates and reactor volume were corrected to reactor temperature. lSiIass spectrometric analysis of the decomposition products showed the presence of the species NOz, NO, N20, and COz. The presence of NO2 and COZ (1) This work was supported by the U. S. Air Force and the Advanced Research Projects Agency under Contract No. AF04(611)9380 and ARPA Order No. 24. (2) J. M. Sullivan and T. J. Houser, Chem. I d . (London), 1057 (1965).
NOTES
3367
was further substantiated by infrared spectrometry. However, attempts to nstablish a mass balance for these species proved unsuccessful.
-0.4
Results
-0.8
The results obtained for the thermal decomposition of TNM over the temperature range 170-223' are presented in Table I. The rate equation for a simple order, single-reactant reaction in a stirred flow reactor is
k,
=
(PO - P)/(P"T)
(1)
where PO and P are the partial pressures, respectively, of the reactant in the gas streams entering and leaving the reactor, r is the average residence time in the reactor, and n is the order of the It is of interest to note that for a first-order reaction in a stirred flow reactor a plot of 1/P os. will yield a straight line, whereas such a plot will be linear only for a second-order reaction in a plug flow or static reactor. Table I: Experimental Kinetic Data for the Pyrolysis of Gaseous Tetranitromethane Run no.
Temp,
TI
O C
sec
17 18 19 20 21 22 23
170.0
1 2
181.2
186.9
3 4 5 6 10 11 12 13 14 27 28 15 16 7 8 9 24 25 26 29 30 31
192.2
198.0 206.7 213.0
215.0 223.2
99.2 169.5 172.9 174.0 123.5 126.1 128.0 14.3 27.0 41.5 43.7 54.7 63.0 13.4 16.9 23.4 24.3 36.8 9.69 9.78 13.7 14.1 5.98 6.70 7.07 4.74 4.89 5.02 2.43 2.50 2.71
Fraction reacted, a
0.171 0.249 0.307 0.283 0.452 0.455 0.479 0.154 0.265 0.355 0.340 0.416 0.468 0.229 0.274 0.310 0.376 0.477 0.266 0.276 0.541 0.546 0.474 0.475 0.510 0.414 0.442 0.410 0.438 0.455 0.417
kl, 8 0 C -1
0.00208 0.00196 0.00256 0.00227 0.00668 0.00662 0.00719 0.0127 0.0134 0.0133 0.0118 0.0130 0.0140 0.0221 0.0224 0.0192 0.0248 0.0248 0.0375 0.0390 0.0856 0.0855 0.151 0.135 0.147 0.149 0.162 0.144 0.320 0.334 0.264
R", 880-1
0.00222
0.00680
-1.2 si M
A -1.6
-2.0
-2.4
-2.8 2.00
2.04
2.08
2.12
2.16
2.20
2.24
2.28
lOOO/T.
Figure 1. Arrhenius plot for the gas phase pyrolysis of tetranitromet.hane.
The first-order nature of the decomposition of TNM is illustrated in Table I, particularly with the data obtained at 186.9'. It can be shown from eq 1 that a plot of a / ( l - a ) vs. 7,where a is the fraction reacted, should be linear for a firsborder reaction. The rate constant, kl, from each experiment is listed in Table I. The best first-order rate constant, klA, at each temperature, obtained from the slope of the best line from an a / ( l - a) vs. T plot, is given in the last column of Table I. The Arrhenius plot (Figure 1) of these rate constants yields the rate expression
0.0130
kl =
0.0232
0.0382
exp( -40,875/RT) sec-'
(2)
The uncertainty in activation energy is about 1.2 kcal/ mole. The possibility of a heterogeneous contribution to the rates obtained in this study has not been completely eliminated. It appears, however, from the high preexponential factor and the magnitude of the activation energy, that the reaction is homogeneous.
0.0856
Discussion
0.144
We believe that the kinetic parameters obtained indicate that the rate-determining step is the rupture of a single C-N bond
0.150
(N02)3C-N02 0.328
+ (No&&*
+ NOz
(3)
(3) W. C. Herndon, J . Chem. Educ., 41, 425 (1964).
(4)A. A. Frost and R. G. Pearson, "Kinetics and Mechanism," John Wiley and Sons, Ino., New York, N. Y., 1956, p 185.
volume 70, Number 10 October 1966
NOTES
3368
This is in agreement with Benson5s6who has suggested that for bond rupture within a molecule a “high” preexponential factor of 1015 to lo‘* sec-l is to be expected. Reaction 3 is then followed by further bond rupture or, more likely, by an isomerization process in which an oxygen atom is transferred to the unsaturated carbon atom. Since the reverse of reaction 3 should have nearly a zero activat,ion energy, the measured activation energy for the pyrolysis of TNM is approximately equal to the dissociation energy of the first C-N bond in TNM if reaction 3 is rate determining. Reaction 3 will be rate determining if the radical formed reacts much more rapidly by another path than by recombination with NOz. The only other apparent mechanism which could give first-order kinetics independent of the buildup of product NO2 would be one in which reaction 3 was rapidly reversible and the rate-determining step involved the reaction
for solid HNE’ fit equally well a simpler mechanism in which the rate-determining step is the rupture of theinitialc-N bond.
The Dissociation of Tetra-n-hexylammonium Iodide in Dichloromethane
by R. A. Matheson
I
NO2
+ (NO2)$.
+products
(4)
I n this case, the measured activation energy would be equal to the dissociation energy of the C-N bond plus the activation energy of reaction 4. Marshall, Borgardt, and Xoble’ have reported kinetic parameters for the condensed-phase pyrolysis of hexanitroethane (HNE). In the case of solid “E, the preexponential factor was 1018.6and the activation energy 38.9 kcal/mole. Within experimental error, this activation energy is the same as that for gaseous TNM. It was postulated for solid HIVE that the ratedetermining step is the simultaneous cleavage of two C-N bonds to give a diradical intermediate.’ This was based partly on t,he high preexponential factor of 10l8.6 which was attributed to the unfreezing in the activated complex of the carbon-nitro system (however, the uncertainty in this preexponential factor is a factor of lo3 to lo4 since the stated average deviation in the activation energy is 10%). Our results with TNM (one carbon atom) indicate that unfreezing of the carbon-nitro system is not required to obtain a high preexponential factor. The mounting evidence that bond rupture yielding two polyatomic radicals is associated with a “high” Arrhenius factor,6,6 supported by unpublished work from this laboratory, makes it apparent that the kinetic results reported (5) S. W. Benson, Ind. Eng. Chem., 56, 18 (1964). (6) S. W. Benson and W. B. DeMore, Ann. Rev. Phys. Chem., 16, 426 (1965). (7) H. P. Marshall, F. G . Borgardt, and P. Noble, Jr., J . Phys. Chem., 69, 25 (1965).
The Journd of Physical Chemistry
Chemistry Department, Victoria University of Wellington, Wellington, New Zealand (Received July 15, 1966)
A recent spectrophotometric investigation’ led to the conclusion that solutions of tetra-n-hexylammonium iodide in dichloromethane contain ion pairs (n-hexyl)F N+I- as well as free ions. Assuming the activity coefficients of both free ions and ion pairs to be unity, an equation2 was derived which related the optical density of a solution of this salt to its concentration. As this equation contained three unknown parameters [the dissociation constant of the ion pair (KD) and the extinction coefficients of free iodide ions (EF) and ion pairs (ep)], it was asserted that measurement of the optical densities of three solutions having different concentrations (Ci) suffices for the evaluation of KD. M for this quantity was obtained from A value of the optical densities of solutions of concentrations 3.86 X lo+, 9.64 X and 4.82 X M. In order to test this assertion, we have examined the consistency of these optical densities with a variety of values KDand activity coefficient assumptions identical with those used in the original work.’ For the values of K D shown in Table I, extinction coefficients EF and eP can be found such that the optical densities calculated from KD, eP, EF, and Ci agree with the experimental optical densities of the three solutions in question, The maximum discrepancy between calculated and experimental optical densities was 1%, which is less than the maximum uncertainty consistent with the manufacturer’s specifications for the instrument used in the measurements. Since none of the extinction coefficients is physically unreasonable, it may be concluded that the measurements reported by Blandamer, Gough, and Symonsl are not sufficient for the evaluation of KD, and that their estimate of this quantity is uncertain by several powers of ten. (1) M. J. Blandamer, T. E. Gough, and M. C. R. Symons, Trans. Faraday SOC.,62, 286 (1966). (2) See eq 3 in ref 1.
