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).
ficient H 3 0 + ion in solution so that the over-all rate is controlled by Na+ ion diffusion. Boggs and Mosher2 studied the effect of HCl pressure only a t a temperature of 385", which was the highest temperature used in their measurements. On the basis of the above reasoning, one can predict that a t a lower temperature where the solubility of HC1 would be greater, there might always be sufficient H30+ ion in the surface layer so that the Na+ ion diffusion would be rate-controlling and the over-all rate would be found to be independI \\ ent of HC1 pressure. We have made such measurements a t 295O, varying the HClpressure between 75 and 430 mm. Within experimental error, the rate of the reaction was found to be independent of pressure under these conditions. The fact that the HC1 curve in Fig. 1 falls very slightly below that for HBr may indicate that even a t 500 mm. pressure of HC1, the H30+ ion concentration on the surface is not sufficiently high to allow complete 1.5 1.6 1.7 1.8 control by the Na+ ion diffusion (evidence for 1/T x 108. which is also found in Fig. 3 in the paper by Boggs Fig. 1.-Arrhenius plot for the reaction of Pyrex #7220 and Mosher2). glass wool with HCI or HBr. H2S and SOz form weak acids when they dissolve in water, giving very low HSO+ ion concentrations. Attempts to measure the rate of the reaction of rate, being controlled by these very low conHBr atl 450" led to irreproducible results. For any The centrations, would be too slow to measure. CHsCl one run, the kinetic results followed the same rate be expected to react by hydrolyzing in the law as a t lower temperatures, but in different runs might surface film of water to form HC1, which would the rate constant varied from 15 X l O - l 9 to 35 X then behave as HC1 alone does. Apparently this 10-19. does not happen, possibly because the hydrolysis Unlike the reaction of HC1 with glass, the rate of reaction removes the surface water before an adethe reaction of HRr with glass was unaffected by quate H 3 0+ ion concentration can be established. the pressure of the hydrogen halide in the gas phase. Several semi-quantitative isotope exchange exVariations of HBr pressure from 30 to 430 mm. periments were performed in an attempt to correcaused no measurable difference in the rate of the late the reaction between the hydrogen halides and reaction. glass with published studies on halogen isotope exWe propose that the reaction of HC1 or HBr with change reactions. Three reaction tubes of the type glass occurs in the following manner, Even a t by Boggs and Brockways were filled with HC1 400" there is a thin film of adsorbed water held te- used 39.7% Cla7 (the normal abundance is containing naciously on the glass surface. The gas dissolves in 24.5% Cy7). tubes were then heated a t 300" this film, forming H 3 0 +ions and halide ions. Pro- for 4 hours to The allow the HC1 to react with the glass tons from the H 3 0 +ions then diffuse into the glass, wall. After this time, were evacuated, filled Na+ ions diffusing outward a t the same rate to with CHaCl containingthey the normal isotopic ratio, maintain electrical neutrality, this ionic counterand heated for 2 hours. The results, sumdiffusion being the rate-determining step in the resealed, marized in Table I, show that exchange of C1 beover-all process, A t 450", the water film is no tween HC1 and CH,Cl can occur through the interlonger maintained intact, and the rate falls off, the mediate formation of NaCl on the wall. exact value of the rate constant being determined TABLE I by the extent to which the water has been removed from the surface. ISOTOPE EXCHANQE BETWEEN NaCl FORMED ON THE REWith sufficient H 3 0 +ions in the surface solution, ACTOR WALLAND CHICl Cl'l in the rate of the counter-diffusion of N a + ions and Pressure Temp., CHsCI, CHaCI, H + ions may be limited by the rate of Na+ difOC. mm. % fusion. Thus above a certain pressure limit, the 150 200 24.6 rate would be independent of gas pressure. At 485 300 25.9 lower pressures of HC1 in the gas phase, the surface 485 450 31.4 concentration of H 3 0 + would be decreased, so that Boggs and Brockwaya found that they could obthe diffusion of H + takes over as the rate-determining step, leading to a slower reaction. It would tain reproducible exchange rates between HC1 and be difficult to predict the relative solubilities of HC1 CH3Cl only if the reaction tubes were evacuated and HBr in a surface film of moisture a t these tem- and heated a t 450-500" for 4 hours before use. peratures, but in bulk solutions a t lower tempera- Otherwise much more rapid exchange was obtures the solubility, on a molar basis, of HBr is served. Since the exchange rates reported in this greater than that of HC1 and decreases less with study are higher than those reported by Boggs and temperature. Thus it might be that over the pres- Brockway in baked tubes, it appears probable that sure range studied with HRr, there is always suf- exchange through NaCl formation was respon,sible
NOTES
Jiine, 19.57
82i
for the highcr rat'es in unbaked t,ubes. The ex- and material (MgCOs) only, hence is thc same for tensive heating served to remove adsorbed water all the differential thermographs, regardless of tli\ from t,he glass surface, so that any exchange mould COz p r e s ~ u r e . ~The constraint must therefore he have to proceed by a different mechanism. The ir- imposed that the rate of reaction is ihe same a t all reproducibility of the measurements in unbaked the reaction temperatures. In order to obtain an tubes could have been caused by variations in the expression for the rate of reaction, we assume the surface area of NaCl exposed on the surface (see rate equat,ion Fig. 1 of Boggs and Mosher2). r = krf(MgCO8) - krPdtsQ(MgO) (3) It is to be expected, of course, that HCl and HBr Here r is the iiet rate of reaction, kr and k, the rate will rcant with glasses of different composition t o quittedifferent extents, depcndent mainly on the so- constants for the forward and reverse reactions, dium content of the glass. The alkali cont>entof respectively, and P d t a the pressure of COz; the Pyrex #i220 glass used in these studies is higher f(MgCOs) and g(Mg0) are functions, respectively, than that of the Pyrex #7740 commonly used in the of the activities of MgCO, and MgO. By defining construction of laboratory apparatus. Pyrex #7740, r as a constant, equation 3 is restricted to those sets however, contains 3.6% Na, and from the results of of temperatures (including the reaction temperathe isotope exchange experiments it appears that it tures) where the rate of reaction is the same. Dividing equation 3 by 1Cr g(Mg0) gives reacts in a similar manner. KINETJC EFFECTS I N DETERMINING HEATS OF REACTION BY DIFFERENTIAL THERMAL ANALYSIS BY HANSJ. BORCHARDT General Engineering Laboratory, General Electric Company, Schenecladu, N e w York Received J a n u a r y 17, 1067
The dynamic gas method of differential thermal analysis as developed by Stone' recently has been used to measure the heat of reaction of the magnesite dissociation. MgCOs I _ MgO t COz (1) Diff erentJial thermographs for this decomposition were obt,ained in atmospheres ranging from 0.001 to 3320 mm. of COz. Increasing CO, pressures displaced the peak to higher temperatures. The In of the pressure was plotted against the reciprocal of the reaction tjemperature3 according to the equation
The term kr/k, is by definition the equilibrium constant which for this reaction is equal tto the equilibrium partial pressure of COz (Peq). Sincc P,, is truly given by the Clausius-Clapeyron equation, P,, = exp( - AH/RT C ) , equation 4 becomes
+
Substituting for k r with the Arrhenius equation, IC = S e x p -E/RT, gives
where S is for all practical purposes a constant and E, the activation energy for the reverse reaction.
Equation 6 relates the reaction temperaturc to the pressure in the DTA sample holder without the assumption made in eqmtion 2 that equilibrium conditions exist. The behavior predicted by equat,ion 6 can best, be seen by the substitution of data. AH is taken as 10.1 kcal./mole, the value given by Stone.2 AH C is evaluated by taking Pe, == 1 atmosphere a t In P d t s = - c' ( 2 ) RT 410" as reported by Cremer and GatL6 The acwhere P d t a is the pressure of COz, T the reaction tivation energy for the forward reaction has been temperature in O K . , AH the heat of reaction, R the observed to be 35.G kcal./mole.6 This gives E, gas constant, and C' a constant. The slope of this as 35.6 - 10.1 = 25.5 kcal./molc. The functions plot being -AH/R gives the desired heat of reac- f(MgC03) and g(Mg0) are, in the absence of data, tion. taken as unity. The constant, r / S is evaluated by This use of a Clausius-Clapeyron type equation taking P d t a = 0.001 mm. a t 350" as reported makes the implicit assumption that the reaction by Stone.2 temperature can be treated as the equilibrium deWith these data, equation G is plotted as a concomposition temperature a t the particular partial tinuous curve in Fig. 1. The circles represent the pressure of COZ employed. The purpose of the data reported by Stone.2 The broken line was present discussion is to examine the validity of this used loy Stone to calculate A H . Relatively good assumption. agreement is obtained. It should be noted howWe first inquire as to the significance of the reac- ever that the agreement hetween the theoretical tion temperature. This i s the temperature at which curve and Stone's data with regard to the limiting the rate of reaction, hence the rate of heat absorption, slope and the temperature a t which P d t a approaches i s sufiiently rapid to establish a temperature dif(4) This assumes t h a t the instrument itself does n o t undergo a n y ferential which the instrument i s just able to detect. significant clinnges in the temperature range under consideration. This minimum detectable rate of heat absorption A formal justification for this interpretation of t h e reaction temperat,ure can be derived from the kinetic equations described by H. J. (rate of reaction) is a property of the apparatus
+
(1) R. L. Stone, J . Am. Ceram. Soc., 86, 76 (1952). (2) R. L. Stone, {bid.. 97 46 (1954). (3) T h e term "rrartion temirernturc" is unrd t o designnt,c tllr tem-
peratrrre at which t h e nenk
iR
first obRerved to anpcar.
Borchardt a n d F. Daniels, J . Am. Chem. Soc., ?e, 4 1 (1857). (5) E. Cremer a n d F. G a t t , Rader Rundaohau, 4, 144 (1949); Ceram. Abrtr., 56d (1950). ( 6 ) H. T. S. Rritton, 8. J. Grcgg and C . W. Winaor, Trans. Faraday Sot., 48, 03 (1952).
