NOTES
356
L
0
25
50 Cd
75 C12Mole%-
I
100
Figure 3. Surface tension isotherms for the (Cd-Cs)Cl system a t various temperatures.
When the size of the alkali metal ion is varied by replacing cesium with rubidium, the shape of the isotherm changes slightly; the minimum becomes less deep, and the maximum is shifted toward higher alkali chloride concentrations. This means that-the alkali chloride concentration being equal-the CdCI8- concentration is lower with Rb+ than with Cs+; in other words, replacing cesium with rubidium causes the complex anion to become less stable. On going from rubidium to potassium, the isotherm changes further in the same way; the maximum and minimum are no longer distinct,, but overlap in a horizontal inflexion a t the composition of an approximately equimolar mixtures. At least, in the (Cd-Na)C1 system, maximum and
minimum have disappeared, and a slight inflexion only is present. We may therefore point out that really the stability of the complex anions decreases gradually from cesium to sodium, in agreement with the considerations outlined in the Introduction. By observing the phase diagrams of these systems,ls we may notice that the larger the alkali metal is, the higher and sharper the peak of the congruent IleCdCL compound. If this may be regarded as a rough indication of the stability of the complex, we clearly find a parallelism between this trend and our previous conclusions. 2. E$ect of Temperature. In Figure 3, the surface tension isotherms of the (Cd-Cs)C1 system a t various temperatures are shown. It appears that with increasing temperature the maximum flattens. We may understand this behavior by making the assumption that pure cadmium chloride contains some associated entities, which are likely to be thermally dissociated, so that their concentration diminishes by increasing temperature. Thus, the dissociating effect of the alkali chloride will be less and less remarkable as the associated entities disappear upon heating. Probably, if it were possible to work at higher temperatures, pure CdClz would no longer contain any associated group, and no positive deviation would appear in the surfacre tension isotherms; these would have the same shape as those of the PbCL systems. (18) E.Degurnov, Dokl. Akad. Nauk SSSR, 64, 517 (1949).
NOTES
Desorption of Cumene from Silica-Alumina Catalysts
by Yutaka Kubokawa and Hisashi RiIiyata Department of Applied Chemistry, University of Osaka Prefecture, Sakai, Osaka, Japan (Received March $8, 1967)
The kinetics of cumene cracking on silica-alumina catalysts have been investigated by many workers. As for the chemisorption of cumene on silica-alumina, there seem to have been no studies made of the heat of adsorption. Only the heat values on catalytically active sites have been estimated approximately from the kinetics of the reaction.' I n a previous work,2 it The Journal of Physical Chemistry
has been shown that desorption rate measurements can give unambiguous information on the heat of adsorption over a wide range for a given chemisorption system. In the present work similar measurements have been carried out for cumene adsorbed on silica-alumina.
Experimental Section Materials. A silica-alumina catalyst containing 13% alumina was obtained from the Shokubaikasei Go. It has a BET surface area of 448 m2/g. Cumene (1) C. D.Prater and R. M.Lago, Advan. Catalysis, 8, 298 (1956); W.B. Horton and R. W. Maatman, J . Catalysis, 3, 113 (1964). (2) Y. Kubokawa, Bull. Chem. SOC.Japan, 33, 546, 550, 555, 739, 747, 936 (1960); J . Phys. Chem., 67, 769 (1963); 69, 2676 (1965);
Y. Kubokawa and 0. Toyama, Bull. Chem. SOC.Japan, 64, 1407 (1962); Y. Kubokawa, S. Takashima, and 0. Toyama, J . P h y s . Chem., 68, 1244 (1964).
357
NOTES was purified by passing it through silica gel heated to 100". Ammonia was obtained from the thermal decomposition of ammonium chloride and purified by fractional distillation. Apparatus and Procedure. A conventional constant volume apparatus was used. I n order to measure the rate of desorption, two traps were attached to the reaction vessel. After the adsorption of cumene at room temperature, the desorption experiment was started by immersing the trap in liquid nitrogen. By using the two traps alternately, the desorption could be continued without interruption. The amounts of cumene collected in the trap were determined by gas chromatography. From the amounts desorbed in a definite time the rates of desorption were determined. During the rate measurements the temperature of the catalyst was lowered abruptly and the rates before the temperature drop were extrapolated to those for the small amount adsorbed after the temperature drop. Thus, the rates at the two temperatures corresponding to the same amount adsorbed and the activation energy of desorption could be obtained. Afterward, the temperature of the catalyst was raised in stages, and similar measurements were carried out at each stage.
Results and Discussion The activation energies of desorption of cumene from a silica-alumina catalyst are shown in Figure 1. It is seen that in the temperature range 20-60' the activation energy of desorption Ed ranges from 12 to 18 kcal/mole, while above 60" it increases markedly with increasing desorption temperature up to 40 kcal/mole. It was found that above 60" the desorption products contained benzene, suggesting that the decomposition of cumene had occurred. According to the work of Horton and Maatman, who investigated the kinetics of cumene cracking on silica-alumina, the heat of adsorption of cumene on active sites is only 9.5 kcal/mole, much less than the E d values shown in Figure l.3 After various amounts of ammonia were adsorbed on silica-alumina, similar additional experiments were carried out, the results of which are again shown in Figure 1. It is seen that the plots of E d against the amount of cumene adsorbed for the catalysts before and after the ammonia adsorption nearly coincide, although no decomposition of cumene took place after the ammonia adsorption of 4.08 cc/g. This suggests that cumene adsorption on silica-alumina is unaffected by the ammonia adsorption; Le., most of the cumene adsorption is not associated with the acidity of silica-alumina. With regard to the cumene adsorption on the active sites for the decomposition, it should be noted that the sites with a high heat of adsorption such as 40 kcal/mole still remain active after the ammonia adsorption. It can be concluded that the sites on which a strong adsorption of cumene occurs are not the same as the
NH, adsorbed 0 ccsrp/gcat
0 0 A
0.548 1.72 4.08
A
.,. A
P 20 .* +
.t
t:
,
4101
,
,
10
,
j
20
Remaining amount of cumene adsorbed, cc(STP)/g of catalyst.
Figure 1. Activation energy of desorption of cumene. Figures indicate the temperature of desorption.
strongly acidic sites, ie., the most active sites for the reaction. I n the temperature range where the decomposition of cumene occurs, the activation energy of decomposition was determined by measuring the benzene content in the decomposition products in a similar manner to the measurements for the activation energy of desorption. The results are shown in Table I. As an ex-
Table I : Activation Energy of Cumene Decomposition
Temp, OC
60 100 150
-Activation Desorption
16.7 24.1 32.9
energy, kcal/moleDecomposition
20.5 32.2 41.9
Remaining amount of cumene adsorbed, cc (STP) / g of catalyst
4.71 0.312 0.033
ample, the rate data used to calculate the activation energy of decomposition are given in Figure 2. The rate shown in this figure was found to be in fair agreement with that expected from the data of cumene cracking a t high temperatures given by Pansing and Mallay.* It is seen that the activation energy of decomposition is increased with increasing temperature, Le., with increasing E d values. With regard to the nature of such an increase in the activation energy of decomposition, further studies are now in progress. At present it seems very difficult to predict whether a similar increase in the activation energy will occur in the (3) In view of a rapid establishment of the adsorption equilibrium, i t may be certain that cumene adsorption on silica-alumina is a nonactivated type. (4) W. F. Pansing and J. B. Melloy, Ind. Eng. Chem., Proc. Design Develop., 4, 181 (1966). Volume 76, Number 1 January 1088
358
NOTEB include all species, paired or unpaired.) Eo' for ferroferri was constant from pH 3.7 to 6.7 and EO'for &-&HZ was constant from pH 2 to 6.7, (These Eo% were defined by E = EO'and E = EO' 2.3(RT/F)pH1 respectively, for the systems in the cited pH ranges, for equal concentration of oxidized and reduced forms. The stoichoimetric equation (1) was confirmed spectrophotometrically by determining the amount of Fe(CN)e3- formed in excess quinone and then in excess Fe(CN)64-.1a The reaction order was investigated by (1) pseudo-first-order plots with ferrocyanide in large excess, and (2) when no reagent was in large excess, by determining best tangents to amount reacted vs. time plots. The slopes and known concentrations yielded apparent second-order k's [ = rate/(ferro).(Q), where s denotes stoichiometric]. All k's were constant a t fixed (ferri),/(ferro), a t pH 2.74 to 4.74. Typical data for I and I1 are plotted in Figures 1 and 2.6 The k's for I11 were constant in this range of ratios. The (ferro), was varied from 0.3 to 4, 0.2 to 2, and 1to 10 ( X M), respectively, and the (Q)%from 0.1 to 1, 1 to 5, and 1 to 10 (X M ) . All data obeyed
+
7 6 5 Remaining cumene adsorbed, cc(STP)/g of catalyst.
Figure 2. The rates of decomposition of cumene at 60". The rate of decomposition a t 40' for the cumene adsorbed of 4.17 cc was found to be 5.64 X 10-6 cc (STP)/g of catalyst, min. Comparison of the rates a t both temperatures corresponding to the same amount adsorbed (4.17 cc) leads to an activation energy of 20.5 kcal/mole.
usual flow method, owing to a marked difference in the experimental conditions, e.g., the reaction temperature or the coverage of cumene during the reaction.
Kinetics of Ferrocyanide Reduction of Quinones
by S. A. Levisonla and R. A. Marcus'b
We describe here the kinetics of ferrocyanide reduction of 2,bdichloro-, benzo-, and 2,5-dimethylquinones (denoted below by I, 11, 111) in 1 M KC1. Reagent materials were used where available and the hydroquinones of I and I11 were synthesized by SnClz reduction and recrystallization. Slightly alcoholic solutions of quinone^^*^ in 1M KC1 were deoxygenated, placed in blackened vessels at 27O, and then in a 10-cm path Beckman cell, Buffered 1 M KC1 ferrocyanide at 27' was added and the rate was measured from ferrocyanide diappearance at 420 mp or disappearance of I at 275 mp, using a B e ~ k m a n . ~ The acid dissociation constants of H2Fe(CN)e2- and HFe(CN)s3- in 1 M KC1 were measured with a Beckman (260-280 mp) at concentrations of 0.4-4.5 X 10-6 M and found to be 0.32 f 0.02 and (4.7 f 0.2) X 10-3 M , respectively'" The equilibrium constant, K , of Q
+ 2Fe(CN)e4- + 2H+ QHz
+ 2Fe(CN)sa-
(1)
was measured from the formal potentials Eo'of the two half-cells using a Leeds and Northrup potentiometer. The Eo' (relative to sce) of the ferrscyanide system in 1M KC1 was 0.22 V, and those for I, 11,and I11 in 1M KC1 were 0.47, 0.44, and 0.34 V yielding K's of (2.8 f 0.5) X 108, (3.1 f 0.5) X 107, and (1.1 f 0.2) X lo4 M-2, respectively. (The anions in (1) are largely paired with K +; concentrations throughout this paper The Journal of Physical Chemistry
C (ferri), (ferro),
1- 1 k kl'
Department of Chemistry, Polytechnic Institute of Brooklyn, Brooklyn, New York (Received July 10, 1967)
(2)
and are summarized in Table I: kl' is fairly insensitive to pH in the range (2.74 to 3.74) where it was measured? C c(: (H+)-' for I, and C a (H+)-'e4 for 11. Table I : Summary of Kinetic Data" -111--
PH
F----I-----(2,5-Dichloroquinone) ki' X C X lo-' 10'
2.74 3.74 4.74 5.74 a
Units for
3.6 3.6
...
0.036 0.39 2.7
--II-
(2,5-
Dimethyl(Benzoquinone)
quinone)
x
cx
10-1
101
ki'
C
3.7 3.7
0.014 0.54 20. 390.
36 26
b b 5
B1'
.,.
,..
kl' and C-1 are ( M min)-'.
b
...
Not measured.'
(1) (a) S. A. Levison, Ph.D. Thesis, Polytechnic Institute of Brooklyn, June 1962. (b) Address correspondence to this author at Noyes Chemical Laboratory, University of Illinois, Urbana, Ill. (2) I and I11 were dissolved in ethanol because of slow dissolution in water and were diluted to 1.0 to 0.01% alcoholic content, a variation without effect on potentials or reaction rates. (3) Versene, lo-' M , added to the ferrocyanide solution to remove a slight turbidity, perhaps due to PbsFe(CN)s, did not affect rates. (4). In the kinetically important time of 10 min, a AT (0.7') was estimated to cause a minor increase of 7% in rate. (6) When (ferri),/(ferro)s changed during reaction in pseudo-firstorder data, a mean was used. In ref la, C and E in Table 32 should be interchanged. (8) A flow system would permit determination of ki' at pH 4.74 and 6.74 by permitting measurements at very low conversions and hence at very low mean (ferri)B/(ferro)eratios, For compound I11 at pH 2.74 and 3.74, only initial rates were measured and no effort was made to determine C. Those initial rates were unchanged when (ferro). was increased from 1 to 10 X 10-4 M .
