initial displacement of the radicals in units of collision diameters of the radicals. PI and Pz are the probabilities of reaction between, respectively, two free radicals and free radical with iodine upon encounter. While we have no independent evaluation of these factors, they are not likely to be much less than unity. Taking PI = Pz = 1 we find y = 1.35 and po = 3.9. These values are larger than those observed in some other s y ~ t e m s ~but . ~ J are consistent with the configuration of the acetyl peroxide molecule and the size of the solvent molecule, since the unit of measurement refers to the methyl radical. (7) J. C. Roy, J. R. Nash, R. R. Williams, Jr., and W. H. Hamill. J . A m . Chem. Soc., 78, 619 (1956).
ADSORPTION OF NITROGEN ON PURE SODIUM CHLORIDE BY DONALD S. MACIVERAND PAULH. EMMETTI Contribution from Multiple Fellowship o/ Gulf Research & Development Company. Mellon Institute, Pittsburgh, Pa. Received J a n w r y 14, 1068
8
Vol. 60
NOTES
824
The Brunauer-Emmett-Teller2 method of determining the surface area of solids by means of isothermal gas adsorption measurements has been found to be applicable to a wide variety of adsorbent-adsorbate systems. Recently, however, Young and Benson3 have reported that the BET plot of nitrogen adsorption at - 195” on very pure sodium chloride exhibits a negative intercept. This effect, if real, would seem to indicate a breakdown of the BET equation when applied t o this system. Therefore, in order to determine the reality of the effect reported by Young and Benson and to explain their results, we have studied the adsorption of nitrogen 4.0
_.3.0
2 X
h
6 -. 2.0 3
5 E40
k 1 .o
0 0
0.1
0.2
0.3
P/Po. Fig. 1.-BET plot of nitrogen adsorption on sodium chloride at -195O. (1) Johns Hopkina University. Baltimore, Md. (2) 9. Brunauer. P. H. Emmett and E. Teller, J . A m . Chsm. ern., 80, 309 (1938).
(3) D. M. Young and G. C. Benaon, private communication.
on sodium chloride at - 195” in the pressure region in which the BET equation should be applicable. Experimental The sodium chloride employed in this work was provided by Dr. G. C. Benson and was prepared by electrostatically recipitating the fumes from molten sodium ~ h l o r i d e . ~ Krior to the adsorption measurements the adsorbent was tramferred in an atmosphere of dry nitrogen from the evacuated, sealed-off bulb in which it was received to a suitable Ram le tube. The voyumetric adsorption apparatus employed in the present work was similar to the one described frequently in the literature .6-7 The nitrogen used for adsorption meaRurement8 was of high purity grade from Linde Air Products Company and was freed of condensable impurities by passage over glam heads at - 195’; the helium, for calibrational purposes, was from the Air Reduction Company an$ was purified by passage over degassed charcoal a t -195 Prior to the adsorption measurements the sample was evacuated at room temperature to a pressure of lo-&mm.; the evacuation was continued until such time as the rate of gas evolution from the sample was negligible. It was necessary to avoid the use of high temperatures in the degassing process in order to prevent any sintering of the sample. A nitrogen adsorption isotherm was then run at -195’. The BET plot of the adsorption data is given in Fig. 1. The straight line obtained up to a relative pressure of about 0.1 had an intercept of 3 X lo-‘ and a slope of 0.105. From these, a “c” value of 351 and a V,,, of 9.49 cc. (STP)/gram was calculated. The latter value corresponds to a surface area of 41.7 m.g/gram, if the adsorbed nitrogen molecule has a cross-sectional area of 16.2
.
Discussion It is obvious from Fig. 1 that the BET plot for this system tends to deviation from linearity a t a fairly low relative pressure. Instead of the customary linear portion extending from relative pressures of 0.05 to 0.35 we find one extendihg from relative pressures of 0.01 to 0.10. Young and Benson made their adsorption measurements at relative pressures of 0.06 to 0.28. As can be seen from Fig. 1 the points from 0.06 to 0.20 lie in a fairly straight line due to the small second derivative of the broken curve. Consequently, if one had only those points measured by Young and Benson one would naturally take the 0.06 to 0.20 region as linear and conclude from this that the plot had a negative intercept. From the curve in Fig. 1, however, one can say that the linear part of the BET plot between relative pressures 0.01 and 0.1 gives a positive intercept and a reasonable surface area. Furthermore, this linear part of the curve extends as it should from a little below to a little above the relative pressure (0.05 in this case) corresponding to a monolayer. It would appear that on certain crystals having smooth faces with very little heterogeneity, adsorption data tend to deviate from the behavior expected, in the direction of tending to give a stepwise adsorption isotherm. This reflects itself in the BET plots by a departure from linearity a t relative pressures that are considerably lower than those a t which the BET plots ordinarily cease to be linear when nitrogen is used as the adsorbate. I n systems of this nature it therefore seems advisable t o extend the adsorption measurements to relative pressures low enough to make sure that one is deal(4) D. M. Young and J. A. Morrison, J . Sci. Inalr., 81, 90 (1954). (5) P. H. Emmett, “Am. Soa. Testing Materials, Symposium on New Methods for Particle Size Determination,” 1941, p. 95. (6) P. H. Emmett, Advances in Coll. Sci., 1, 1 (1942). (7) P. H. Emmett and 9. Brunauer, J . A m . Chem. SOC.,66, 35
(1934).
1
NOTES
June, 1956 ing with the linear portion of the BET plot that includes the relative pressure at which a statistical monolayer is obtained. THE DECOMPOSITION OF MALONIC ACID IN GLYCEROL AND I N DIMETHYL SULFOXIDE B Y LOUIS WATTS CLARK Conlribslion from the Department of Chemiatry, Saint Joseph Collsge. Emmilaburg, Maruland Received January 18. 1066
A number of investigators have made kinetic
studies on the thermal decomposition of malonic acid alone and in solution.' Although it is well known that malonic acid, like oxalic acid, is smoothly decarboxylated in glycerol, the kinetics of the reaction in that solvent have not been previously reported. Preliminary studies in this Laboratory revealed the interesting fact that the solvent, dimethyl sulfoxide, likewise promotes the decomposition of malonic acid, and does so even more efficiently than glycerol. The kinetics of the decomposition of malonic acid in these two solvents have been carefully studied in this Laboratory, and results of this investigation are reported herein. Experimental Rea ents.-C.P. malonic acid was further purified by recrystafiisation from ether The purity of the reagent was demonstrated by the fact that the volume of carbon dioxide evolved from every quantitative sample in the decarboxylation experiments was invariably stoichiometric. Dimethyl Sulfoxide (99.9% ure), and glycerol, Analytical Reagent Grade, 95% by vofume, were also used in these experiments. Apparatus and Technique.-The apparatus and technique used in these experiments have been previously described.2 Decomposition of Malonic Acid in Glycerol.-At the beginning of each experiment 100 ml. of glycerol was placed in the dry reaction flask in the thermostated oil-bath. A 0.1671- sample of malonic acid (sufficient t o eld 36.0 ml. of 80,a t STP on complete reaction) was pcced in a thin glass capsule (blown from 6 mm. soft glass tubing and weighing approximately 0.29.) and introduced at the proper moment into the solvent in the manner previously described. The rapidly rotating mercury seal stirrer immediately crushed the capsule, the contents were dissolved and mixed in the solvent, and reaction began. The evolved carbon dioxide was measured at constant pressure. The above procedure was repeated at eight different temperatures between 150-160". Every sample of malonic acid yielded the stoichiometric volume of carbon dioxide within ex erimental error. For example, the final observed volume ofgas at STP produced by one sample at 154.3' was 35.6 ml.; a t 156.0", 36.0 ml.; a t 156.3', 35.9 ml.; at 159.9', 36.0 ml.; at 161.0°,36.0 ml. Duplicate and triplicate runs a t the same temperature showed excellent reproducibility. Decomposition of Malonic Acid in Dimethyl Sulfoxide.At the beginning of each experiment 100 ml. of dimethyl sulfoxide was saturated with dry carbon dioxide gas and laced in the reaction flask in the thermostated oil-bath. gamples of malonic acid weighing 0.1671 g. were added to (1) (a) J. Laskin, Trans. Sib. Acad. Agr. Po?., 6, No. 1 (1928); C.A., 13, 1804 (1928); (b) J. Bigeleisen and L. Friedman, J . Chem. Phus., 17, 998 (1941); ( 0 ) G. A. Hall, Jr., J . Am. Chsm. Soc., 71, 2891 (1949); (d) J. G. Lindsey, A. N. Bournes and H. G. Thode, Can. J . Chsm., 19, 192 (1951); (e) J. G. Lindsey, A. N. Bournes and H. G. Thode, ibid., 80, 183 (1952). (2) (a) H. N. Barham and L. W. Clark, J . Am. Chsm. Soc., 73, 4638 (1951): (b) L. W. Clark, ibid., 77, 3130 (1955); (c) 77, 6191 (1955).
825
the solvent and the evolved gas messured aa described above. The experiment waa repeated at five dflerent temperatures between 130-140'. Every sample invariably yielded the stoichiometric volume of carbon dioxide allomng for the experimental error. For example, the final observed volume at STP produced by one sample at 129.8" wan 36.1 ml.; at 134.3 , 35.5 ml.; at 138.8 , 36.0 ml. Reproducibility was excellent.
Results and Discussion The experimental data were converted to standard conditions and milliliters'of evolved gm was plotted against time for each temperature. Values of x corresponding to different values of t were obtained from the resulting isotherms. Log (a x) was then plotted against t (a is the theoretical stoichiometric volume of carbon dioxide, 36.0 ml.). The points thus obtained for the middle 80% of the reaction fell on perfectly straight lines in every experiment. This fact indicates that the decomposition of malonic acid in glycerol, m well as in dimethyl sulfoxide, is a first-order reaction. From the slopes of the lines thus obtained the specific reaction velocity constants for the decomposition of malonic acid in the two solvents were calculated for the various temperatures. For the cme of the decomposition of malonic acid in glycerol, the temperatures studied, as well as the corresponding speciiic reaction velocity constants in sec.-l, were M follows: 152.2")0.00364; 154.1°, 0.00413; 155.7",0.00470;156.3', 0.00483; 158.1", 0.00513; 159.0",0.00570; 159.9",0.00578; 161.0°, 0.00666. Results for the case of the decomposition of malonic acid in dimethyl sulfoxide were as follows: 129.8", 0.00416; 131.3", 0.00490; 134.3", 0.00618; 137.0°,0.00740; 138.8",0.00845. A straight line was obtained in each case when log k was plotted against 1/T according to the Arrhenius equation. From the slopes of the lines thus obtained the energy of activation and the frequency factor for the reaction in each solvent was calculated. For the decomposition of malonic acid in glycerol these were found t o be 25,500 cal., and 5.3 X lo"', respectively; for the decomposition of malonic acid in dimethyl sulfoxide corresponding values were 23,350 cal., and 2.3 X lO*O, respectively. The temperature coefficient for the reaction in glycerol was found to be 2.11,in dimethyl sulfoxide 2.26. The enthalpy of activation, entropy of activation and free energy of activation at 140°,according to the Eyring equation, were found to be as follom: for the decomposition of malonic acid in glycerol, 24,600 Gal., -12.2 e.u., and 29,650 cal., respectively; for the decomposition of malonic acid in dimethyl sulfoxide, 22,300 cal., -15.0 e.u., and 28,350cal., respectively. It is of interest to compare the rate of reaction in the two solvents a t some particular temperature. At 140",k for malonic acid in glycerol is 0.0013,for malonic acid in dimethyl sulfoxide it is 0.0092. Therefore, malonic acid decomposes seven times aa fast in dimethyl sulfoxide as it does in glycerol at this temperature. It is also of interest to ascertain how the values of k in these two solvents compare with that for the pure acid for a given temperature. By melting one-half mole (53 g.) of the pure acid
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