ROBERT L. ALTMAX
366
to have a much higher electron attachment coefficient than oxygen. 6. Summary A. sensitive ionization chamber electron swarm method has been applied to measurement of the electron attachment of a number of gases a t atmospheric pressure. With some exceptions, a qualitative correlation was found between electron attachment and
Vol. 67
flame inhibition effectiveness, i.e., the higher the electron attachment coefficient, the greater the effectiveness of the inhibitor. The effective halogenated hydrocarbon inhibiting agents were found to exhibit a higher electron attachment coefficient than oxygen. However, from other work it is clear that a high electron attachment coefficient alone is not indicative of effectiveness as an inhibiting agent.
VAPORIZATIOK OF MAGNESIUM OXIDE AKD ITS REACTIOX WITH ALUMINA BY ROBERT L. ALTMAN* Lawrence Radiation Laboratory, University of California, Lit ermore, California Received August 17, 1963
A Knudsen effusion and oxygen transpiration study of the vaporization of MgO indicates that gaseous RIgO is of little importance in the vapor phase a t about 20OO0K. A Doo for gaseous MgO of about 80 kcal./mole is obtained. Knudsen experiments with MgAlzOt support a negative heat of formation from the oxides, and this is confirmed by recalculation of other experimental results.
Introduction A mass spectroscopic study of the vaporization of MgO a t 1950°K. has shown the predominant ion peak to be Mg+ with NgOf less than 1/1000 the Mg+ intensity. The MgO dissociation energy obtained with these data is much lower than that obtained from Knudsen effusion weight loss determinations,2 or from magnesium flame intensity rneasurement~.~-~ I n an attempt to resolve these conflicting results, further Knudsen effusion and transpiration experiments with MgO and with MgA1201have been performed. Experimental The general characteristics of effusion and transpiration techniques are described elsewhere.6 The experimental apparatus for the Knudsen cell studies included an Ajax spark-gap generator to heat by induction a suitable crucible and its MgO or hfgA1204 contents. Magnesium oxide crystals of 99.7% purity obtained from the Norton Company were used in all the AIgO Knudsen and oxygen transpiration experiments. A small quantity of powdered MgA1204 was supplied by E. G. King of the U. S. Bureau of Mines (Berkeley). Its preparation and properties are described elsewhere.? Tungsten crucibles had successfully been used to heat A1208 in Knudsen effusion experiments,8 but Brewer2 and Ackermanng report that such crucibles react with MgO to produce a volatile tungsten oxide. Therefore, the ready availability of Morganite1o alumina crucibles and previous successful heating of MgO in an alumina liner1 prompted its use in the present work. Magnesium aluminate was formed by reaction between the MgO and the alumina container in both the Knudsen and transpiration runs. Much of the MgO gaseous species condensed on the crucible walls as MgA1204. The Knudsen effusion work using pure MgA1204as the starting material was performed in order to -
test whether the total weight loss from the MgO runs was affected by the formation of magnesium aluminate. Temperatures were determined by sighting a Leeds and Northrup optical pyrometer upon the effusion orifice in the lid of the crucible. The pyrometer was calibrated against a Sational Bureau of Standards calibrated pyrometer by comparison of the temperatures read by the two instruments when sighted on a tungsten lamp. The initial MgO effusion runs were made with alumina crucibles and lids used just as received from the manufacturer. The weight-loss results obtained were greater than could be accounted for by assuming vaporization to the elements. HOPIever, further experiments indicated that vapor leakage through openings other than the effusion orifice occurred. Therefore, both the alumina crucible and the lid were reground to obtain a better fit. A tungsten cap was made to fit over the alumina lid and upper part of the crucible in order to reduce the possibility of leakage through the lid. Experiments with closed crucibles modified in this manner showed only the expected weight loss due to the vaporization of A1203 in a tungsten environment.1l n’eight loss data have been obtained with three different sets of alumina crucibles and effusion orifices. Two crucible lids had their effusion holes ground to knife edges in order to eliminate the need for a Clausing correction. The third lid had a cylindrical orifice and when this crucible was used, the appropriate Clausing correction (0.3) was applied t o the geometric hole area. The partial pressures were obtained from the total weight loss, W , by application of the Knudsen equation
W (g./hr./cm.2) = (1.596 X lO5/Ti”)
C PiMii’z
( i
)
where Pi are the partial pressures of the effusing species and M , are their molecular weights. Assuming that MgO vaporizes chiefly by decomposition to the elements, we have
* Alameda Strtte College, Hayward, California. (1) R. F. Porter, W. A. Chupka, and M. G. Inqhram, J. Chem. Phys., 29, 1647 (1955). ( 2 ) L. Brewer and R. F. Porter, zbzd., 22, 1867 (1954). (3) E. M. Bulewioz a n d T. hI. Sugden. Trans. Faraday Soc., 56, 7200 (1959). (4) L. Huldt and A. Lagerqvist. Arkzv Fyszb, 2, 333 (1950). ( 5 ) I. V. Veits and L. V. Gurvich, Optiba z Spektroskopzya, 1, 22 (1956). (6) J. O’M. Bockris, J. L. White, and J. D. Mackensie, “Physioochemieal Measurements a t Bigh Temperatures,” Butterworth’s Scientific Publications, London, 1959, pp. 225-246. (7) E. G. King, J . Phys. Chem., 69, 218 (1955). ( 8 ) L. Brewer and A. W. Searcy, J . A m . Chem. SOC.,73, 5308 (1951). (9) R OJ. Ackermann and R. J. Thorn, ”Reactions Yielding Volatile Oxides a t High Temperatuw,” XVI-th International Congress of Pure and -4pplied Chemistry, Paris Long Islnnd City, N. Y.
and
c PiMi’/z i
+
= P3fgfl~nrg1’~
+
POXO”~
POzMOz1/2 (4)
(11) J. Drowart, G. DeMaria, R * PABurns, and 31, G , Inghrami J 8 Chemi Z’hys,, 32, 1366 (19SO)i
Feb., 1963
~ A P O R I Z A T I O KO F hIAGNESI‘tJM O X I D E AKD ITS
REACTION W I T H ALUMINA
367
According to eq. 1, the number of moles of material effusing is proportional to the partial pressure divided by the square root of its molecular weight. Therefore
t
I 0-41
Equation 5 substituted into eq. 4 yields
0
/
0
/
/
= 8.18Pxg and with eq. 1
(7) The points shown in Fig. 1 were obtained by application of eq. 7 to the experimental weight loss data. It has been observed that the weight loss of empty alumina crucibles in similar experiments is of nearly the same magnitude as crucibles containing MgO. Hence, the experimental partial pressures can be expected to be high by perhaps a factor of two. The dashed line in Fig. 1indicates the theoretical PJfeexpected from the appropriate thermodynamic data. It was obtained by substituting eq. 2 and 3 in eq. 5 , Le.
and the resulting cubic equation solved by standard methodn.12 The equilibrium constants, K1 and K z ,were computed from tables of free energy functions and appropriate heats of formation.l:I Figure 1 also gives the PJI, resuIts of Brewer and Porter.2 Our weight loss results indicate that gaseous MgO is of much less importance than found by Brewer and Porter. This can be shown more clearly by computing the heat of vaporization with eq. 2 for each experimental point and comparing the value with the thermochemical value of 178 kcal./mole. By eliminating eq. 3 from the P Mcalculations, ~ we obtain the results given in Table I. TABLEI KNUDSEN VAPOR PRESSURE OF MgO ‘/*Oz(g) W X s ) ----f Mg(g)
+
T
(OK.)
1884 1976 2035 2007 2120 2099 2013
Orifice area, cm.2
8.86 X ( X 0.3)
AHao
PYg ( a h )
2.79 x 10-6 1.83 x 2.20 x 10-6 9.68 x 2.43 x lo-; 2.35 X 3.72 x
(koal./mole)
166 176 169 172 176 174 178 173
5.42 3.19 3.52 8.49 5.30 7.25 1.33
2009 1947 1978 2065 2039 2018 2044
x
x x
x x
x x
10-5 10-5
10-6 10-5
175 173 175 178 178 176 173
2.3l X
2.43 8.82 1.28
x
x x
10-7 10-7 10-5
0
I 11
t 5.8
5.6
5.4
I x T
I
I
I
I
5.2
5.0
4.0
4.6
I
4.4
io4( o K - ~ ) ,
Fig. 1.-Knudsen vapor pressure of MgO: squares, 8.56 X ( X 0.3 Clausing correction) cm.2 orifice area; triangles, orifice area; filled circles, 2.31 X 8.48 X 10-3 (knife edge) 10-2 (knife edge) cm.2 orifice area; open circles, Brewer and Por Iter. least 2 kcal./mole. It would seem, therefore, that the difference between the experimental heats and the calculated thermochemical value are within the experimental uncertainties of the pressure measurements. In further support of the Knudsen results, experiments have been performed in which oxygen gas has been passed over crystalline MgO a t temperatures above 2000’K. and the weight loss has been interpreted as gaseous MgO. These results are given in Table 11. From the weight loss data, the heat of the reaction MgO(s) = MgO(g) obtained by the third law method was calculated to be 158 dz 4 kcal./mole. The measured weight losises were near the Emit of sensitivity in measurement and may be looked upon more properly as yielding a lower limit for the heat of sublimation of MgO. It can readily be shown that a heat of sublimation of 158 5 4 kcal./mole will increase the weight loss in a Knudsen effusion experiment by an insignificant amount above that due to gaseous Mg and oxygen.
-
TABLEI1 MgO TRANSPIRATION RESULTS WITH OXYGENAT 1 ATM. MgO(s) MgO(g)
175 1747 1916 1999
E 4..
169 178 170 172
Because these results also include the weight loss of the alumina crucible, the heats calculated from these data are too low by at (12) “Handbook of Chem. and Phys.,” Chem. Rubber Publ. Co., Cleveland, Ohio, 41nt Ed., 1960,p. 318. (13) G. N. Lewis and M. Randall, “Thermodynamics,” revised b y K. .:f Pitzer and 1,. Brewer, McGraw-Hill Book Go,, Inc., New York, pi, Y., 2nd Ed., 1961, pp. 669-686.
2’ (OK.) 2033 2140 2150 2145 2165 2154 2165 2175
Moles/ hr.
t (hr.)
20.0 14.9 16.7 16.6 16.5 16.6 16.6 16.5
4 4 6 5 6 6 7 7
Wt. loss (mg.)
AHaO
P M ~ (atm.) O
3 . 0 x lo-* 3.3 X 1 . 7 x 10-7 .2 6.0 x .I 2.5‘x .3 7.5 X .6 1.3 X 0.3 6.4 x
0.1 .8 .7
(kcal./rnole)
155 152 ’ 156 160 ’ 165 160 158 162 158 zk 4
The free energy function data used for gaseous MgO assumed a ‘I:ground state. If, however, there is a low-lying ~IIstate as in CZ,l4the free energy function around 2000°K. will be larger (14) L. Brewer and
s, Trajmar, J . Chem. Phys.,
86, 1586 (1’962).
ROBERTL. ALTMAX
368 I
I
I
I
l
/
'
I
I
d
/I
//
o -t~ L
j
c
/
5.6
1
I
1
I
I
5.4
5.2
5.0
4.8
4.6
T
x IO~(OK-~),
Fig. 2.-Knudsen vapor pressure of Mg-41204: filled circles, cmeaorifice area. I is calculated assuming a heat 2.31 X A1203) of - 5 kcal./mole, of formation from the oxides (MgO and I1 is calculated assuming a AHr of zero.
+
by R In 6 and the heat of sublimation will be increased by 7 kcal./mole. The dissociation energy for MgO gas can be estimated by subtracting the results of Table I1 from the atomization energy of MgO, 237 kcal./mole. I n this manner, an upper limit of 80 kcal./mole can be obtained which will only decrease somewhat if the ground state is other than lx. This dissociation energy value is in good agreement with the results of a BirgeSponer extrapolation16 and this apparent agreement will be the subject of a subsequent pap+. Both the Knudsen effusion and transpiration work have shown that the volatility of MgO around 2000°K. is almost entirely due to elemental species. Therefore, a comparison of the Knudsen volatility of MgA1204 with our MgO results should allow an estimation of the heat of formation of magnesium aluminate. Kalyanram and Bell16 and Schma1zriedl7report a free energy of formation for this spinel from MgO and A1203 of -5 t o -6 kcal./ mole at about 1500%. Combining this value with appropriate free energy function data13 leads to a heat of formation from the oxides of about -5 kcal./mole a t 298°K. This may be compared to an experimental value of +1 kcal./mole reported by Grjotheim, Herstad, and Toguri.18 The Knudsen results shown in Fig. 2 were obtained in the same manner as described for MgO. Comparison of the MgAL04 magnesium partial pressures with those of MgO given in Fig. 1 indicates that the formation of bIgA1204 by heating MgO in alumina crucibles does not make the magnesium partial pressure equal to that obtained by heating pure PlllgAh04. Chemical z,nalysis of the residue showed that the alumina content of the spinel increased upon heating. However, X-ray powder patterns failed t o reveal any of the characteristic lines of AlzOs. It appears that the alumina produced upon heating MgAlzOl dissolved ih the remaining magnesium aluminate, further decreasing the activity of the magnesium vapor. Magnesium aluminate is known to incorporate excess alumina in this temperature range.1Q It wm found that the lowest vapor pressure results a t the higher temperatures contained the greatest excess of alumina in the residue. If the decomposition of magnesium aluminate is written as
(15) (1943)
4. Lagerqvist, Arkiv Matematik, Astronomi och F y s i k , ZQA, 25
(16) hl. R . Kalyanram and H. B. Bell, Trans. Brit. Ceram. SOC., 60, 135 (1961). (17) H. Schmalzried. Z. physik. Chem., 26, 178 (1980). (18) K. Grjotheim, 0. Herstad, and J. M. Toguri, Can,. J. Chem., S9, 443 (1961). (191 D. hl. Roy, R. Roy, and E. F. Osborn, A m . J . Boi., 261, 337 (1953).
T'ol. 67
third law calculations of the heat of this reaction from the observed magnesium vapor pressure will be somewhat in error unless the reduced activity of the alumina is included in such calculatiors. This error will be leest in those runs in which the activity of the alumina is greatest. As shown in Fig. 2, somewhat better agreement between experimental acd calculated magnesium partial pressures, particularly at the highest temperatures, is obtained with a negative heat of formation of magnesicm aluminate from the respective oxides.
Discussion Porter2 obtained a maximum value for Pntg0 a t 195OOK. of 7.5 X lo-* atm. He further observed that the ratio, I+M,/I +M,O was greater than 1000 and that I +&I, greatly exceeded 1+02. Our oxygen transpiration results at slightly higher temperatures yield a dissociation energy for llIg0 about 10 kcal./mole less than Porter's value of 90 kcal./mole. Under neutral conditions the magnesium partial pressure in equilibrium with MgO(s) a t l!XO0K. is atom.20 Because the ionization cross section of Mg is about five times that of oxygen,21the P M , computed atm., larger by a from Porter's findings is at least factor of 100. Ackermann and Thorn20 suggest that the MgO(s) was subjected to reduction by the surrounding tantalum heater. We find that if these equilibria, are assumed
+ nfgO(s) = hIg(g) + TaO(g) = TaO(g) + O(g) TaO(g) + l/202(g) + ~ l g o ( g >= TaO(g) + Rfgk)
Ta(s) Ta(s) TaOdg) TaOdg) Tab)
+ 2MgO(s) = 2Mg(g) + Ta02(g) =
and a consistent set of thermodynamic data applied to ~ 2 - ~pressures ~ the solution of these e q u a t i o n ~ , ~ , ~ 3 8these are obtained Mg
1.0 X atm. 5.9 x 10-9 02 1.8 X lo-'" TaO 1.2 X 10-6 TaOz 1.5 X
o
These pressures taken together with the MgO pressure obtained by Porter yield an MgO heat of sublimation of 145 kcal./mole and a dissociation energy of 90 kcal./ mole. Since these results are not significantly different from those obtained by considering the sole reaction to be RfgO(s)
=
MgO(g)
it is felt that the reductant effect of Ta upon the 3lgO pressure is negligible. Therefore, the heat differences between our transpiration results and the mass spectrographic data are believed to be due to experimental error. Our Knudsen effusion work with MgA1204 seemingly favors a negative heat of formation for this spinel from MgO and &03. In order to obtain a more precise heat value, we have applied the third law to the magnesium partial pressures reported by Grj otheim, et al.,'s for the reaction 4NIgO(s)
+ 2-41(1) -+- 1LIgL21204(s)+ 3Rlg(g)
(20) R. J. Ackermann and R. J. Thorn, Progr. Ceram. Sci.,1, 48 (1961). (21) J. W. Otvos and D. P. Stevenson, J . A m . Chem. SOO., 78, 546 (1956). (22) L. Brewer and M. S. Chandrasekhsriah, UCRL-8713 (June, 1960). (23) L. Brewer and G. M. Rosenblatt, UCRL-9437 (October, 1960). (24) M. G. Inghiam, W. A. Chupka, and J. Berkowitz, J . Chem. Phys.. 27, 669 (1967).
CHARACTERIZATION OF ADSORPTION RATESON HETEROGEXEOUS SURFACES
Feb., 1963
Because magnesium is soluble in liquid aluminum,25 the usual expression m0208 --
TABLE I11 HEATOF REACTION FOR 4MgO(s) f 2Al(l) + MgA1204(s)
A(FTO - H m O )
T
T
- R I n -p a M g U2A1
(OK.)
for calculating the heat should include the reduction in A1 activity produced by the solution of magnesium. This can be done by solving the equationP
X M-~ (1600/RT)(X~1)~
In
aMg
= In
In
aAl
= In XAI
-
T
(1600/RT)(X~,)~
to obtain the composition and activity of aluminum in equilibrium with magnesium vapor. The results are given in Table 111. Omission of the activity correction would increase the heat by 0.5 to 1.5 kcal./mole. The lower heat value, 127 kcal./rnole, yields a heat of formation for MgA1,04 from the oxides of - 6 kcal./mole and for the upper limit of 131 kcal., a heat of formlation of -10 kcal./mole. Other than possible errors in the values of the free energy functions used in these calculations, no explanation for the decreasing heat value is apparent. Ifowever, these results lead to an (25) A. Sohneider and E. K. Stoll, Z. EEektrochem., 47, 519 (1941). (26) “Selected Values for the Thermodynamic Properties of Metals and Alloys,” Minerals Research Laboratory, Univ. of Calif., Berkeley, Calif ., 1958.
Pnaga (mm.)
+ 3Mg(g) - A(FTO H%ss)/TC
a
d
XM,~
1143 3 . 5 5 0.047 0.085 1194 7.56 .OS5 .095 1232 13.3 .065 .lo9 1271 23.6 .078 .117 1345 56.0 .095 .146 .114 .170 1379 89.0 .119 ,177 1388 100.1 1390 102.5 .120 .177 .139 .200 1414 143.2 Reference 26 a Reference 18.
aAib
369
(e.u.)
aH0298
(koal./ mole)
0.910 .900 .884 .877 .843 .816
82.94 130.9 82.47 130.8 82.20 130.4 81.94 129.8 81.48 129.6 81.26 128.6 .808 81.10 128.2 .808 81.10 128.1 .782 81.07 127.3 References 7, 13, and 27.
estimate of the heat of formation of magnesium aluminate from the elements of -552 kcal./mole, but ithe error limits of this value are difficult to assess. Acknowledgments.-This work was performed under the auspices of the U. S. Atomic Energy Commission at the Department of Mineral Technologyof the University of California, Berkeley, and LRL-Livermore. The author is indebted to Professor Alan W. Searcy of .the Department of Mineral Technology for his guidance a,nd support during this investigation. (27) K. R. Bonnickson, J . Phys. Chem., 69, 220 (1955).
A NOVEL TECHNIQUE FOR CHARACTERIZATIOS OF ADSORPTION RATES ON HETEROGENEOUS SURFACES BY L. M. NAPHTALI AND L. M. POLINSKI Polytechnic Institute of Brooklyn, Brooklyn 1, hr. Y . Received August 10,1962 A novel technique is proposed for obtaining and interpreting data on adsorption rates to a catalyst surface. The method is illustrated by actual data from a hydrogen-on-nickel system. The amount of adsorbed gas on a catalyst which is part of an isothermal system varies with time when the pressure changes. The variation depends on the adsorption kinetics and the hetero,geneity of the surface. For a sinusoidally varying pressure, the dependence of the adsorption amplitude and phase lag on the frequency is one way of characterizing adsorption kinetics. The “frequency response” to an induced sinusoidal pressure variation of the moles of gas adsorbed on a uniform surface having first-order kinetics can be computed theoretically. A heterogeneous surface is assumed t o be an assembly or series of “different uniform-surfaces” randomly interspersed. An assembly of such surfaces, characterized by different rate conetants, has an out-of-phase component of the adsorption which resembles a spectrogram, separating the effect of different types of surface sites irrespective of the fact that adsorptions are occurring simultaneously on &like sites. As an illustration of the technique, hydrogen adsorption was studied on a supported nickel catalyst. The effect of oxygen addition t o the catalyst on the adsorption kinetics of hydrogen was studied. It was found that an increase in oxygen content reduced the amount of fast adsorption and increased the slow adsorption. It was possible t o characterize and separate the rates of adsorption of both the fast and the slow types.
Introduction The problem to be discussed is, essentially, how to interpret data on adsorption rates to a catalyst surface. Frequently during chemical adsorption on a catalyst surface, several processes occur simultaneously. This may be due to the heterogeneous nature of the surface or to the existence of different adsorbed states. I[n either case, the tools of automatic control theory can be used to separate the phenomena and give further information on their nature. I n order to determine the dynamic characteristics of an unknown system, ,the control engineer uses or induces certain forms of disturbances of “inputs” and observles or interprets their effects or “outputs.” Two of the
most useful types of inputs for the study of process dynamics are the step-function and the sine wave. I n the study of the kinetics of adsorption, the input commonly used (though not usually thought of in this sense) is the step-function, or instantaneous jump in a system variable. A typical example of a step input is an evacuated chamber containing catalyst which is instantaneously opened allowing the adsorbing gas to enter freely. The sine wave input, however, has been overlooked in studies of adsorption kinetics.1 The sinusoidal disturbance is particularly useful when the surface being investigated is expected to be hetei~o(1) L. Polinski, Doctoral Dissertation, Polytechnic Institute of Brooklyn, 1961, p. 29.
