August, 1958
SOLUBILITY OF PENTAERYTHRITOL TETRANITRATE
1009
SOLUBILITY OF PENTAERYTHRITOL TETRANITRATE' BY ROBERT N. ROBERTS AND ROBERT H. DINEGAR University of California, Los Alamos Scientijk Laboratory, Los Alamos, New Mexico R e c e i v e d May 14, 1068
The solubility of pentserythritol tetrsnitrate (PETN) in acetone, aqueous acetone, ethyl alcohol, ethyl acetate and benzene has been measured as a function of temperature. The data have been analyzed in terms of the e uations of solutions relating the tem erature coefficient of solubility with the heat of solution of the solute. The heat of !usion of PETN has been estimated Lorn these experiments as 23-24 kcal. mole.-'
Introduction Pentaerythritol tetranitrate, PETN, is an explosive organic nitrate ester that has been known for more than fifty years. It is normally prepared by the simple nitration of pentaerythritol, during which process the four hydroxyl groups are replaced by nitrate groups.2 The preparation of pentaerythritol, however, by the condensation of formaldehyde and acetalde hyde, in the presence of Ca(OH)2,has been shown by Fredrick and Bruna to give rise to certain byproducts, the most important being dipentaerythritol. This latter consists of two pentaerythritol nuclei, formed from two of the methylol groups produced by the reaction of three moles of formaldehyde with one mole of acetaldehyde and the subsequent reduction of the aldehyde group of trimethylol acetaldehyde by a fourth mole of formaldehyde, joined by an ether linkage. Dipentaerythritol, DIPE, gives a hexsnitrate upon nitration, DIPEHN, which must be separated from PETN by its solubility difference in aqueous acetone. a It is thus evident that for purification reasons alone precise solubility data on PETN are desirable. An investigation of the available literature brings to light that this problem has been investigated at least thrice bef0re.~.6 Since these data, however, are not in agreement with each other, a general reinvestigation of the problem for certain common, useful solvents was attempted. Experimental Procedure Excess dried PETN, three times recrystallized from acetone, (m.p. 140.1°),6was added to the solvent' in a solution flask connected by an immersion filter to a weighing bottle. The entire apparatus was then thermostated and the solution stirred magnetically until equilibrium was attained.8 After the given time a portion of the solution was transferred to the weighing bottle by means of a hand as irator. After removing the entire apparatus from the tfermostat, the receiving flask was detached and the weight of PETN per wei h t of solvent was found in the following manner. 8eighing of the sample ffask enabled the weight of solute plus solvent to be determined. Distilled water8 was then (1) Thia work was done under the auspices of the Atomic Energy Commission. (2) L. Vignon and F. Gerin. Compt. Tend., 188, 690 (1901). (3) W. Fredrick and W. Brun. Be?., 63, 2681 (1930). (4) T. Urbanski and B. Kwiatkowski, Roczniki Chem., 18, 741 (1933). (6) J. Tranchant, Mcm. P o u d r e s , 84, 117 (1962); P. Aubertein, Mem. P O U d T e 8 , 84, 107 (1962). (6) Original commercial product was obtained from Trojan Powder Company, Allentown, Pa. (7) Analytical reagent grade, used without further purification. (8)I t was found t h a t after several hours equilibrium was est,ablished. All solutions mentioned in this report were held at constant temperature for at least 8 hours before analysis of the Bolution took plaoe.
added to the flask and the entire sample heated to constant weight a t about 9 0 O . 1 0 The weight of solid PETN was then determined and the solubility easily calculated.
Discussion of Experimental Results The solubility of PETN in g. PETN/kg. solvent, for the solvents acetone, aqueous acetone, ethyl acetate, benzene and ethyl alcohol, is given in Tables I and 11. TABLE I Solubility g. PETN/&. solvent
Temp., A.
Acetone 288.2 293.2 298.2 303.2 313.2 323.2 293.2 303.2 313.2 323.2 293.2 303.2 313.2 323.2 333.2 293.2 303.2 313.2 323.2 333.2
208.10f0.10 248.40 f 0.50 305.60f0.50 345.30f0.20 449.20 f 0.20 587.603t3.70 Ethyl acetate 106.19f2.07 140.62 f 0.15 185.00f0.23 241.93f0.43 Benzene 2 75f0.15 4.96 f 0.08 8 . 3 4 f 0.18 14.48f0.31 25.89f0.03 Ethyl alcohol 1.25 0.01 2.13 f 0.01 3 . 7 8 f 0.02 6 . 5 7 i 0.18 11.96It0.06
It easily can be demonstrated that a plot of the natural logarithm of the mole fraction of solute dissolved (In Nz) versus the reciprocal of the absolute temperature is a straight line for all the solvents investigated. Their slopes, the temperature coefficients of solubility, differ, however, and are shown in Table 111. In general, they are observed t o decrease with an increase in solubility.'l (9) The addition of water is neoessary whether it be miscible with the aolvent (aoetone, ethyl alcohol, ethyl acetate) or not (benzene) to prohibit the occlusion of solvent as the sample approaches dryness. (10) Drying temperatures above 90° cannot be used due t o decomposition of solid PETN. (11) The chance that a change in molecular weight of the dissolved P E T N contributed to the various values of the slopes obtained waa obviated by measuring the apparent molecular weight by the method of Lhe elevation of the aolvent boiling'point. The value obtained for P E T N in all four solvents was 318 zt 1-a figure less than 1% different from the monomeric PETN value of 316. It is evident that no appreciable assooiation, dkaociation, etc., oocura in these solvents.
ROBERT N, ROBERTS AND ROBERT H, DINEGAR
1010
TABLE I1 Aqueous acetone % wat,e; by wt.
Temp., A.
293.2
Solubility,
g. PETN/kg.
solvent
6.23 12.30 18.22 23.99 35.11 55.80 6.23 12.30 18.22 23.99 35.11 55.80 6.23 12,30 18.22 23.99 35.11 55.80 6.23 12.30 18.22 23.99 35.11 55.80
298.2
303.2
313.2
162.91f0.17 93.12f0.05 52.22f0.10 28.68f0.05 6.80 f 0.03 0.25 f 0.004 19O.86fO0.09 1 1 4 . 6 2 f 0.10 66.12f0.08 36.66f0.05 9.25 f 0.03 0.38f0.003 225.92 f 0.26 137.37 f 0 . 2 1 81.83f0.10 46.53f0.04 12.57 f 0 . 0 3 0.57f0.009 314.21 f 0.03. 202.53f0.17 126.58f0.26 76.60f0.19 23.30f0.17 1.26 f 0.01
Vol. 62
with fa = activity coefficient of the solute in the saturated solution. If one considers the relationship between the temperature dependence of the Gibbs free energy change and the thermodynamic equilibrium constant, the van’t Hoff isochore results In-K -d =
AH0
(4) dT RT’ where AHo = difference in the molal heat content between the standard state of reactants (pure solute H) and products (solution of infinite dilution, HO). Substitution of K = fZNzin (4)gives
Since Equation 5 becomes
Nowla Substitution in (6) gives
TABLE I11 A In
Nz
Solvent
cal. mole-1
Ethyl acetate Acetone Benzene Ethyl alcohol
4300 4970 10945 11000
and, for any concentration of solute
Consider the following equation which represents the solubility equilibrium solid solute (SZ) = saturated solution (SSz)
(1)
The thermodynamic equilibrium constant K in terms of activities ( A ) ,is AssnlAsn
(2)
Since the activity (Asz) of the pure solute is unity, by definition K
In general AH8,,1
TABLE IV A In N2
fzNz 1
wt. % water
0 6.23 12.30 18.22 23.99 35.11 55.80
190.00 30’ 40 50 60 70 80 Wt. 70HzO. Fig, 1.
+ AH,
AHf = AHf*,the calorimetric (true) heat of fusion only when there is no interaction between solute and solvent-and at this point AH, = 0, making,
(3)
I
10 20
AHi
AHf = heat of fusion of solute (at the specific conditions) AHm = heat effect of mixing liquid solute with solvent
or K
=
with
90
4.97 5.89 6.99 8.10 8.97 11.36 15.06
x
108
23.4 AHt*
simultaneously AHsOl = AHf”. Thus [d In Nz/ d(l/T)],,t is proportional to aMf*, a constant, only (12) G. N. Lev& and M. Randall, “Thermodyasmica,” McGrawBill Book OQ., Now York, W. Y,, 1823, p. W0.
V
August, 1958
REACTION OF HYDROGEN ATOMSWITH SOLIDOLEFINS
(1) when the solution is ideal, or (2) when the solution is at infinite dilution. Since none of the PETN solutions are ideal the organic solvent experimental results are in line with equation 8 which predicts a lower temperature coefficient of solubilitv in those solvents that have the greater disso1ving"power as indicated by [b In fi/ bNz]~. The PETN aqueous acetone solutions likewise are non-ideal and it is seen in Table IV that R [A In Nz/
I011
A(l/T)ISatincreases as the weight per cent. water in the aqueous acetone mixture increases-the solution becoming correspondingly more dilute at saturation with respect to PETN. In addition to an interpretation similar to that mentioned for organic solvents. the aaueous solutions allow an esh a t i o n of A k f * , f o r an extrapolation of R [ A In N2/A(l/T)]sat to pure water (infinitely dilute PETN solution) meets criterion (2) of equation 8. Table IV and Fig. 1 show these results.
THE REACTION OF HYDROGEN ATOMS WITH SOLID OLEFINS AT -19501 BY RALPHK L E I NAND ~ ~ MILTOND. S C H E E R ~ ~ National Bureau of Standards, Washington, D.C. Received May 16, 1068
Hydrogen atoms, produced on the surface of a hot tungsten ribbon, react readily with certain olefins condensed a t - 195'. Of the olefins investigated, propene, butene-1, 3-methylbutene-1, 2-methylbutene-1 and isobutene show rapid hydrogen uptake. Butadiene-1-3 pentene-1 and 3,3-dimethylbutene-l react very slowly. trans-Butene-2 and hexene-1 show no measurable reaction. +hese rate differences are ascribed to small differences in activation energies which readily are observed at these low tem eratures. The mechanism of the reaction appears to involve two consecutive H atom additions in which an alkyl radicaf resulting from the hydrogenation of the olefin, can be further hydrogenated to the alkane. The possibility of alkyl radical stabilization in sizeable concentrations is indicated.
I. Introduction Hydrogen atoms react with some solid olefins at 195°.3 The occurrence of this reaction provides a possible method for the preparation of alkyl radicals stabilized in a matrix. A study of the kinetics of the low temperature hydrogen atom addition process gives detailed information concerning the relative reactivities of various olefins. It is found that hydrogen atoms diffuse readily through solid hydrocarbons and, in the case of olefins, reaction can occur through the bulk of the solid. The existence of measurable rates of reaction under these conditions makes possible the detection of small differences in activation energy. Propene, for example, can be hydrogenated almost completely, while hexene-1 does not react. Extensive work has been reported on the gas phase H atom addition to ~ l e h s , ~ -but - ' ~considerable variation exists in reported values for the activation energy, Melville and Robb give the value of 2.5 kcal. per mole with little difference among the olefins studied. Rabinowitch, Davis and Winkler give 10.8 kcal./mole for propene, but this is not consistent with the occurrence of the addition re-
-
action at -195". Darwent and Roberts give the value 5 kcal./mole for D atom addition to propene. Analysis of the dimerization products for this reaction shows that addition occurs almost completely on the terminal carbon of a primary ~ l e f i n . Our ~ observations are in accord with this work.
II. Experimental
The addition of H atoms to solid olefins a t -195' was observed by exposing the olefin, uniformly condensed on the inner surface of a one-liter bulb, to atoms produced on a heated tungsten ribbon.8 The ribbon, 2 cm. by '/z cm., was welded to heavy nickel leads. These in turn were mounted in the vessel through kovar-glass seals. The metal components were electropolished prior to final assembly. The ribbon was heated electrically and the current was controlled by an auto-transformer operating from a constant voltage source. The temperature of the filament was observed with a micro-optical pyrometer. A layer of high purity olefin, lO-6-lO-* cm. thick, was deposited on the inner surface of the reaction vessel. Hydrogen was introduced by diffusion through a heated palladium thimble to pressures of from 20 to 100 p . Pressures were measured by a thermocouple gauge, overdriven for increased sensitivity, and calibrated for hydrogen with a McLeod gauge. Changes of as little as 0.1 p could be measured. Rates of reaction were followed by recording the preRsure drop, pressure measurements being made with the tungsten ribbon a t operating temperature. (1) This researoh wag performed under the National Bureau of Observations have been made on the hydrogen pickup by Standards Free Radicals Research Program, supported by the De- the following olefins, listed in order of decreasing rate: propartment of the Army. butene-1, isobutene, 3-methylbutene-1, Z-methyl(2) (a) Guest Scientist, Oh-Mathieson Chemioal Corporation; utene-1, pentene-1, 3,3-dimethylbutene-1, butadiene-1,3. (b) Guest Scientist, General Electrio Co., Cincinnati, Ohio. The last three of these have very slow rates. Hexene-1 and (3) R. Klein and M. D. Scheer, J . Am. Chem. Soc., 80, 1007 (1958). trans-butene-2 show no measurable pickup. (4) K. H.Geib and P. Harteck, Ber., 66B, 1816 (1933). The analysis of the products of the hydrogenation of (5) W. J. Moore, Jr., and H. S. Taylor, J . Chem. Phys., 8 , 604 butene-1 and propene, using a mass spectrometer, gave the (1940). results shown in Table I. The rate a t which atoms impinge on the olefin surface is (6) D. J. LeRoy and E. W. R. Steacie, ibid., 10, 676 (1942); 10, 683 (1942). proportional to the rate a t which they are formed by dis(7) H. W. Melville and J. C. Robb, Proc. R o y . SOC.(London),8196, sociation of H z on the hot ribbon. This in turn is governed 494 (1949). by the rate of arrival of hydrogen molecules a t the ribbon (8) B. de B. Darwent and R. Roberts, Dis. Faraday SOC.,14, 65 (determined by the pressure) and the ribbon temperature. (1953). The rate of hydrogen uptake by butene-1 as a function of the (9) W. J. Moore, Jr., and L. A. Wall, J . Chem. P h y s . , 17, 1325 tungsten ribbon temperature was measured. The rates (1949). were taken as the reciprocal of the time required for the (10) B. S. Rabinowitch. 8. G. Davis and C. A. Winkler, Can. J . pressure to decrease from 31 to 25 p . The results are Rse.. B21, 251 (1943). shown as an Arrhenius plot in Fig. 1.
gene,
