NOTES
824
Vol. 01
ionic surface-active agents should not neglect the induction of ion-selective properties in the membranes by adsorption of the surface-active ions.
AN ACTIVE SPECIES FORMED I N T H E E1,ECTRICAL DECOMPOSITION OF DIRlETHYLAMINE BY FRANCIS OWENRICEA N D CHESTER GRELECKI Contribution from the Chemistry Department, Calholic Univsreily o/ America, Washinpion, D . C. Received February 8, 1967
t2
t
t
t c,/c*,
-
Fig. l.-A, the dependence of Emsx(-) and Emem( -) on CI/CZfor dodecyltrimethylammonium chloride when each of the concentrations is at or above the critical micelle conand Emem centration; B, the de endence of Ern.=(-) (- -) on CI/CI for &decyltrirnethylammonium chloride when each of the concentrations is a t or below the critical micelle concentration; C, the dependence of E,,, (-) and E m a m (-) on CI/C* for potassium chloride.
-
-
selectivity of the membrane when an appreciable quantity of the surfactant is in micellar form. This is because (1) adsorption of the dodecyltrimethylammonium ion is maximal and (2) it is much more difficult for the micelles, because of their large size and high positive charge, to migratc through the membrane than it is for unassociated surface-active ions. The data given in Table I1 and shown in C of Fig. 1 show that the membrane itself has little or no ion-selective character when placed between solutions of KCI. The electromotive force observed is opposite in sign to that expected if the membrane were inherently more permeable to negative ions than positive ions. The positive values observed for E,,, may indicate that the membrane has a slight electronegative character due to the presence of a few acid groups in its structure, although included in Emem is the sum of all of the liquid junction potentials in the cell. I n the light of the above results it appears that a cellophane membrane, as used in equilibrium dialysis experiments with ionic surfactants, can become ion-selective in such a manner as to tend to prevent the passage of surface-active ions or their micelles through the membrane. Presumably this is due to the adsorption of the surface-active ions on the membrane. The precise size of the effect depends on the ratio of the concentrations of the surfactant solutions and on the absolute magnitude of the concentrations. The effect is more important when micelles are present in the system. It appears, therefore, that any explanation for the low rates of equilibration in dialysis experiments with
L.
/,,*-I_-
The decomposition of dimethylamine has been studied thermally2 and photolytically* and the results of these experiments indicate that the main reaction is different in the two cases. Carter, et u Z . , ~ have reported a kinetic study of the methylamines; their analytical results indicate that in an unpacked bulb, the products of the thermal decomposition of dimethylamine consist mainly of methane and a non-volatile residue having the empirical formula CzHsN corresponding to methylmethyleneimine (CHIN=CH2). The results of the photolysis of dimethylamine have been reported by Bamford.8 The actual experimental results showed that the main products of the photolysis of gaseous dimethylamine were hydrogen and a polymer which analyzed (CaHaN).. In our present experiments, dimethylamine was decomposed in the field of a high frequency oscillator and the products were condensed as a green solid a t 196". This material slowly changes to a white solid when allowed to stand a t liquid nitrogen temperatures for extended periods and this color change is accompanied by the evolution of hydrogen containing small amounts of methane. We tried t o prepare the green material by quickly cooling the products of the thermal decomposition of dimethylamine' but were not successful. We were unsuccessful in attempts to prepare the green solid by the photochemical decomposition of dimethylamine.
-
Experimental Anh drous dimethylamine vapor waa passed through the fie& of a high frequency oscillator. With the pressure of the dimethylamine between 10-1 and lo-* mm. a glow discharge resulted when the point of a commercial type Tesla coil was placed on the tube. The vapors leaving the field were quickly condensed onto a surface cooled with li uid nitrogen. The condensate thus formed was a green sol8 a t -196". The permanent gases formed during the initial phase of the reaction were removed from the system by means of a high speed mercury diffusion pump and then analyzed on a mass spectrometer. Anal. Found: He, 94.5,95.0; CHa, 4.5,3.8; Nz,1.0,1.2. The reen deposit slowly turned white even at -190"; when alfowed to warm up a t the rate of about 5" per minute, a sharp transition occurred when the temperature reached about 173". At this point the deposit turned white rather suddenly, there was some blistering and a permanent gas wa8 evolved. The composition of the permanent gas was: HZ, 82.4,81.5; CH4 16,3, 17.6; Np, 1.190.7. Half-life Determinations.-In our early experiments it was observed that the end of the Tesla coil had to be near (1) Thie work was supported in part by the Atomic Energy Commission, Contraot Number AT-(40-1)-1305. (2) A. 0.Carter, P. A. Boaanquet, C . G. Silcocks, N . Travera and A. F. Shilahire, J . Chem. Soc., 495 (1939). (3) C . H. Bamford, ibid., 17 (1939). (4) See F. 0. Rice and C. Grelecki, J . Am. Chem. Soc., 79, 2679 (1957).
NOTES
Julie, 1957 the colt1 siirface otherwise no grccn material was produced. In ordrr to ilt?t'crmine thc half-life of the active spccies, a serics of cspei~iinai~ts were performed in which the distance of the coil from tho cold finger was systcmaticnlly varied and tho tiinc for thc initial appearance of the green compound was iiotcd. This time was taken to bc inversely proportiond t,o tlie conccntrcttion of the active spccies. We ulso mwsrirrd, for various distances of the coil from the fingcr, thc ainouiit of I)ermanont gaR liberated during the transition of tlie green material and assumed that this is direc:t,ly proportional to the concentration of radicals stabili d . 111 thcsc esperimcnt,s the cross sectional area of the tube was 0.785 cln T l ~ Ijressure e drop along the tube was 0.035 mm. per o111.nnd t,hp :tverago prnssure at thereaction sitewas 0.07 niin. Sincc 5.42 X lo-' mole of dimethylamine was pnsficd t,liroiigh the tube in 15 minutcs the average flow rate W ~ 200 R m./sec. s z
TABLE I D is t,hc d i h n c e of the coil from the cold finger. 1 is the time t.ii.ken for the radicds t,o flow from the point of origin to tho liqiiid Nz coolcd siirirtce and is calculated by dividing the tlistanre I,y the flow r d c . p is the pressure of permanent gas formed nftcr t8he tramition of the green materid. I,,, is the h i e rcqiiired for R visible green deposit t o form. D , cui.
1.5 2 3 4
t X 104, sec.
P, mm.
tm, 8ec.
0.75 1 .00 1.50 2.0
0.70 .20 .07 .03
30 240 480
Thc first,-order rate constant was obtained b y Plotting 1% P At ; L r l d a h from the plot l / t m At* T h e rate constan" 'vas the same (1.87 x 10' set*-') regardless of which rnct,hocl was used t,o follow the reaction. The halflife of tho grrrn mat,crial in the vapor phase is therefore 3.7 X lo-' sCc* Tllifi is an extrcmeb short life compared t o the half-lives of other radicals which have been re orted. It is smaller, 1)y a factor of about 25, than the hayf-life6 of N H and thc hydrocarbon radical^^,^ CHa and CH?, and b y a factor of ahout 250, than the half-life of the hydraeino radical.8 The Activation Energy of the Reaction.-When the green solid was allowed to stand for extended eriods at -196' i t became apprccinbly li hter in color. 8ince evolution of hydrogen acconipanied &is reaction, the increase iii pressure could lie used to measure the rate of the reaction at the tempcratiire of boiling nitrogen and boiling oxygen: The temperature of the coolant was measured t o f 0 . 2 with a low tcmperature thermometer calibrated by the National Bureau of Standards. Care was taken t o ensure t h a t the conrcntration of the green compound was the same at the beginning of each reaction, so that the initial rates could be comparcd by means of the Arrhenius equation t o yield an energy of activation for the reaction. The initiul rate of appearance of permanent gas at -195,30 1.42 lo-a mm, per min, at - 1 8 3 . 3 0 it was 4.6 x 1 0 - 3 min. per min. which gives an activation energy of 1.3 krnl. This result is probably accurate t o within a few tenths of a kilocnlorie since an error of 1' in the measure of the t.cmpernture difference results only in an error of 0.1 koal. in the activation energy.
825
assuming that one of the six hydrogen atoms attached to carbon is knocked off in the electric field CH3NHCH2 is and the on the cold finger.
-
T H E REACTION OF ACID GASES WITH
PYREX GLASS1 BY JAMES E. ROQOS, J,YNDAL. RYANAND
TAAURIGL L. PEEK
Department o/ Chemistry, The Uniuarsily of Tezas, Austin Z#, Tezas Received February 83, 1967
In an attempt t o understand the r'nechanism of certain halogen isotope exchange reactions, Boggs and Mosher2have recently studied the chemical reaction between Pyrex $17220 glass wool and HCI gas* They found an extensive, diffusion-contro11ed reaction forming NaCl on the glass StlrfaCe, Proceeding a t a measurable rate in the temperature range between 295 and 385". It is the purpose of the present study to attempt to discover what other gases react in a similar manner, to determine the mechanism of the reaction, and to relate the reaction with Pyrex glass to the effect of the wall in halogen isotope exchange reactions. Experimental IIydrogen bromide was proparod by distillation from an acetic acid solution purchased from East,man Kodak C o , The other gases used Were purrhnsed in cylinder4 from Tho Mathcson c0. ~ 1 gases' 1 urified by distillation in a Rystcm. $he apparatus and exper;mental methods in this study were similar t o those reported
wcrL
fra&,n~
Discussion We have not been able t o determine the cornposition of the green material, but its extremely short life tinle and its reactivity a t very low temPerat?res @ronglY suggcst that it is a free radical stablllixd In the condensed Phase. w e Were not able to detect any trace of tetramethylhydrasine despite the fact that we examined the ~ L S specS the Products very carefully. At the Present t ~ - ~of m time we think that our results arc best explained by
Results and Discussion Hydrogen bromide wa9 found to with pyrex #7220 glass woo1 in 8 Inmner very similar to that previously reported2 for HCI. Electron microscope photographs of the glass surfacc after reacshowed the Of thicknesses Of loosely-adhering crystalline material, which when studied by electron diffraction techniques proved to be NaRr, The specific reaction const'ants were determined by the method of Boggs and Mosher12using an average of 48 experimental points a t each temperature, these being taken from several separate runs using different quantities of glass wool. The shape of the curves and the reproducibility of the individual observations were similar to the earlier work. At H B pressure ~ of 50 mm., the rate constant, k , has a value of 3.7 X 10-lg (moles cm.-2)2 sec.-l a t 300°, 12 X 10--lga t 350" and 32 X lO-l9 a t 400". An Arrhenius-type plot of these values, shown in Fig. 1, gives an activation energy for the reaction of 17 kcal./mole. Figure 1 also shows the results obtained by Boggs and Mosher2 for the reaction of HC1 with glass wool at a pressure of 500 mm. The two curves are nearly the same, the difference being barely more than the limits of experimental error. Under similar conditions, the gases H2S, SO2 and CH3C1did not react with Pyrex #7220 glass WOO] at, a measurable rate.
(5) F. 0. Rice and M. J. Fresmo, J . A m . Chem. Soc., 1 5 , 5529 (1951). (6) F. 0 . Rice and W. R . Johnston, ibid., 66, 214 (1934). (7) F. 0. Riae and A . 8. Glasebrook. ibid.. 6 6 , 4329 (1938). (8) F. 0. R i m and F. Scherber. ibsd., 17, 201 (1855).
(1) This work was supported by grant 4478 from The Wniversify of Texas Research Institute. (2) J. E. Bogas and H. P. Mosher, J . A m . Chem. Soc., 7 8 , 3001 (1956). (3) J. E. Bog@ and L. 0. Brockway, ibid., 11, 3444 (1955).
