2279
C O M M U N I C A T I O N S T O THE EDITOR
Comment on
66
Evidence for Nitrogen
Trioxide in the Combustion of a
+
Double-Base Propellant"
Si,.: In studies on the combustion of double-base propellants by rapid-scan mass spectrometry, Dauerman, Salser, and Tajimal have claimed evidence for the production of NO, as a major intermediate product in the preignition stage when the surface temperature is 200". Since the observed nz/e 46/30 ratio was inverse to that of NOz, another species giving both wz/e 30 and 46 was present. Nitroglycerine, vaporized from the propellant, was excluded as a possible source because it was said to have a cracking pattern totally unlike that found; in particular, it had an nz/e 47 as large as 46 as well as numerous other strong peaks. It was proposed that the species responsible for these observations was XO3. This is not in accord with results obtained here2 on the cracking pattern of nitroglycerine using an A.E.I. Al.S.2H mass spectrometer with a direct inlet system operating at 100" and the source at 200". The values in Table I are the only observed peaks of height 1% or more of the base peak (46).
Table I d e
Height
16 26 27
2 1 1 6 15 24 5
28 29 30 31
product of propellant combustion at low pressures would seem improbable. Thus Guillory and John&ona give R, = atm-l at 25" and AH" = -7 i 2 kcal for the system Oz NO z 00x0; hence [KO] >> [NO,], which is not in accord with Dauerman, Salser, and Tajima's observations. It is also difficult to propose any plausible mechanism of formation of NOS; C-0 bond fission in nitroglycerine is highly improbable since D C - O N O ~ 57 kcal mole-1 as compared with D O - N O ~ 37-40 kcal mole-'. Measurement of the initial rates of evolution of NOz from a range of liquid polynitrates4 gives activation energies of 37-40 lrcal mole-'; similar results mere obtained by Levy,6 who measured the rate of disappearance of ethyl nitrate in vapor-phase pyrolysis experiments.
m/e
42 43 44 45 46 76
- -
(1) L. Dauerman, G. E. Salser, and Y . A . Tajima, 69, 3668 (1965); A I A A J., 5, 1501 (1967). (2) R. T. M. Frnser and N. Paul, J . Chem. Soc., in press. (3) W. A. Guillory and H. S. Johnston, J . Chem. Phys., 42, 2457 (1965). (4) L. Phillips, Nature, 160, 753 (1947); Ph.D. Thesis, London, 1949. (5) J. B. Levy, J . Amer. Chem. SOC.,76, 3254 (1954).
EXPLOSIVES RESEARCH AND DEVELOPMEKT ESTABLISHMENT MINISTRY O F TECHNOLOGY WALTHAM ABBEY,ESSEX,U. K. RECEIVED ~IARC 11,H1968
L. PHILLIPS
Height
1 5 3 1 100 9
It is difficult to envisage any plausible major fragmentation process which gives m / e 47; the spectra of a wide range of mono- and polynitrates examined here2 show no m / e 47. The n z / e 46/30 ratio appears to be consistent with Dauerman, Salser, and Tajima's results, but they report no m / e 76; in their instrument the nitroglycerine peak at this latter nz/e value may be less than that recorded above. This suggests that the hypothesis of formation of NO, is questionable. NO3 is certainly a known species,, but is considered to be a complex much weaker than either Nz04 or N203,so that its detection as a major
Activation Energy for the Mobility
of the Hydrated Electron
S~Y:To obtain an estimate of this activation energy, use is made of the recent suggestion that the mobility of the hydrated electron, eaq-, arises from its movement between existing potential traps.l These are formed by the random orientation of neighboring water molecules. When they are occupied by the electron, apparently no or only little orientational polarization occurs. To account for this, it is assumed that the electron changes its trap within a time equal to or shorter than the dipolar relaxation time in water, namely 10-1' sec. The average displacement, L, of eaQ-can be estimated from the relationship I; = 4207, where D denotes the diffusion coefficient of eaq- and r the dipolar relaxation time. If we take for D = 4.5 X cmz (I) D. C. Walker, Quart. Rea. (London), 21, 79 (1967).
Volume 72, Number 6 June 1968
COMMUNICATIONS TO THE EDITOR
2280 Electric field, $ [ V
that there is some restriction on the movement of the electron on the top of the energy barrier, so that the entropy of activation, AS *, is negative. For this case, an upper limit of the activation energy, E,,,, can be estimated by using the known equivalent conductance2 of eaq- (AeBq- = 177 ohm-’ em2)in eq 2. E,,, was found to be 2.3 kcal mole-’. It may be of interest to conjecture that if an eaq- reaction requires no energy in excess of that for the transfer of eaq-, its activation energy should be close to the E m a x value of 2.3 kcal mole-’. I n agreement with this, a recent measurement of the activation energy for the eaqNOz- reaction gave the value of only 1.7 kcal mole-’. 5 Acknozcledgnzent. The author wishes to thank Dr. XI. Ebert for his helpful criticisms and discussions during the writing of this note.
cm-1).
T
+
Distance, om.
Figure 1 . Free eiiergy profile for the transfer of eaq-
sec-1 and 7 = lo-” sec, L is calculated to be 3 X em, a reasonable value if compared with the average intermolecular separation in water of about 2.9 X lo-* ~111.~
The free energy profile for the transfer of eaq- can be diagrammatically represented as shown in Figure 1. From the barrier height, AF *, and its decrease in the presence of an external electric field, 6F*, the upper limit of the activation energy for the transfer of eaq- can be estimated. The method employed was to calculate the barrier height by applying the theory of absolute reaction rates to the equivalent conductance of eaq-. The derivation is analogous to that for the equivalent conductance of the hydrogen and only the final equation for the equivalent conductance of eaq-, Aeaq-, is given by
where 8’1 denotes the Faraday constant (96,500 C), F Zis the factor to convert electron volts into calories (23,060 cal eV-’), I and L are distances in cm as indicated in Figure 1,($is the potential gradient of 1 V crn-l, R is the gas constant, h is Planck’s constant, 7c is the Boltzmann constant, AF’ is the free energy of activation, and T is the absolute temperature. Introduction of the numerical values into eq 1 and substitution of the well-known thermodynamic relationship
+
(2)
Of the The Overall entropy change for the electron is zero. However, it is reasonable to assume
T h e Journal of Physical Chemistry
LABOR.ITORIES CHRISTIEHOSPITAL AND HOLTR A D I U M I N S T I T U T E AIANCHESTER 20, EXGLAND
B. CERCEIC
ACCEPTED A N D TRANSMITTED BY THEFARADAY SOCIETY (DECEMBER 1, 1967)
On the Critical and Pseudocritical Pressure of Binary Mixtures‘
Siy: This communication may be regarded as a supplement to the paper, recently published, where the relations between the critical constants To,Pc, V c , and the pseudocritical constants Tcz,Pc,, Vcz of mixtures were properly tested except the relations between the pressures P c and PC,. It was supposed that either of these relations is not valid for large differences in critical temperatures of the components, TCl and T c z ,e.g., for the CZRP, n-C7H16system. The error results from a graphical method employed for evaluation of the derivative (b In P,/b In T)v,, appearing in the relation3
+
( b In P,/d In T)v,(TC- Tc,)/Tc, (1)
*
for AF yields the following equation for the activation energy, E , of the eaq- transfer 2.3 log 7TAS* ea4 ,,0°>
(5) B. Cercek and M. Ebert, J . Phys. Chem., 72, 766 (1968). PATEHSON
(PC- PC,)/PCZ=
AF* = E - RT - TAS*
1
(2) K. H. Schmidt and W. L. Buck, Science, 151, 70 (1966). (3) J. Morgan and B. E. Warren, J . Chem. Phys., 6 , 666 (1932). (4) S. Glasstone, K. J . Laidler, and H. Eyring, “The Theory of Rate Processes,” McGraw-Hill Book Co., Inc., New York, N. Y., 1941, pp 563-565.
Present calculations, carried out by means of a computer for 24 systems, show that eq 1 (as well as the (1) Partial support of the American Petroleum Institute is gratefully acknowledged. (2) A. Kreglewski, J . Phys. Chem., 71, 2860 (1967). (3) J. S. Rowlinson, “Liquids and Liquid Mixtures,” Butterworth and CO.,Ltd., London, 1959.
