NOTES
Jan., 1963 ing is continued for 6-12 hr.; no further carbon monoxide is recovered after 6-8 hr. The chemisorption is defined as the difference between gas admitted and gas recovered. The H2-D2 exchange rates also are measured after reduction for 12 hr. a t 482'. The catalyst is cooled to liquid nitrogen teniperature in hydrogen, after which a mixture of 25% D2 and 75% HZis passed over the catalyst a t a constant rate and a t a constant pressure of 518 mm.; the effluent is analyzed for HD by maiw spectrometry. From several consecutive runs a t varying flow rates, the exchange rate constant is attained as the slope OF a linear plot of -In (1 X/X,) us. l / S V , where X is the conversion to HD, X , the equilibrium conversion, and S'v is expressed as cc. (STP)/g./min.
-
Results and Discussion A commercial catalyst consisting of 0.6% Pt om alumina, partially deactivated by use and subsequent regenerations, was treated with dry air at various temperatures. The changes in carbon monoxide chemisorption and platinum solubility are shown in Table I. Sample A, the regenerated catalyst, was found to have a small amount of platinum metal, barely detectable by X-ray diffraction; this plus the low carbon monoxide chemisorption value indicates that some growth of platinum crystallites has occurred. Correspondingly, less than half of the platinum was soluble, compared to nearly 100% for the fresh catalyst. As the severity of oxidative treatment was increased, the extent of platinum solubility increased. At the same time, the platinum surface area upon Subsequent reduction incrertsed. Thus, the formation of platinum-, alumina complex results in greater platinum surface area, as measured by carbon monoxide chemisorption. The existence of a platinum-alumina complex is, therefore, desirable, not per se, but because it leads to a high dispersion of the platinum upon subsequent reduction. Consider the reaction
Pt
+ Oa J_ PtOz
PtOz.alumina
It is clear that in the presence of an oxygen partial pressure, there will be a tendency for the oxide to form, and that the oxidized state will be stabilized by complex formation with alumina. Therefore, an increase in oxidation severity increases the extent of complex formation. At the same time, since the particles of platinum are dispersed by the complex formation, subsequent reduction produces more highly dispersed platinum metal. There are limits to the improvements one can attain by increasing oxidation severity. As platinum-alumina is treated in a given oxygen partial pressure at successively higher temperatures, a temperature will be reached at which the oxygen pressure equals the decomposition pressure. Above this temperature, there will be decomposition to form platinum metal. The critical tempera1,ure wiIl increase as oxygen partial pressure increases. If the reasonable assumption be made that mobility is greatest in the oxidized state, the treatment at a temperature such that the decomposition pressure exceeds the ambient pressure will result not only in the formation of platinum metal, but in an increase in crystallite size. Herrmann, et uL19 have reported a decrease in the amount of soluble platinum on heating in air a t 593 and 625'. The results given in Table I1 illustrate this hypoth(9) R. A. IIcrrmann, S. F. Adler, M. S. Goldstem, and R. R.1. DeBaun, J . Phus. Chem., 6 6 , 2189 (1961).
201
esis. As oxygen partial pressure is increased with temperature held constant, the amount of platinumalumina complex increases. As in Table I, this increase is accompanied by a subsequent increase in platinum dispersion upon reduction, as measured by the rate of H2-D2 exchange. Also, as temperature is increased at constant oxygen pressure, the extent of complex formation passes through a maximum, as predicted. This critical temperature, at which the oxygen partial pressure equals the decomposition pressure, is near 510' at 0.21 atm. and near 580' a t 1.O atm. It is concluded, therefore, that in the oxidized state a platinum-alumina complex does exist, but that it is converted to metal on treatment with hydrogen. Furthermore, the degree of dispersion of this metal increases with an increase in the fraction of platinum soluble in the oxidized state. Acknowledgment.-The authors wish to acknowledge the assistance of J. S. Melik in determining rates of H2-D2 exchange.
T H E KINETICS OF T H E OXIDATION OF HYDROGEN PEROXIDE BY CERIUM(1V) BYG. CZAPSIU,B. H. J. BIELSKI, AND N. SUTIN Chsmistry Department, Bronkhaven National Laboratory, Upron, Long Island, New York neC&ved M a y B8, 1063
Baxendale has reported that the oxidation of hydrogen peroxide by cerium(1V) is complete within a few seconds at room temperature, even at micromolar concentrations of the two reactants, and has proposed that the reaction proceeds in two one-electron steps.2 The kinetics of the hydrogen peroxide-induced cerium(II1)-cerium(1V) exchange in 0.8 N sulfuric acid solutions have been studied by Sigler and Masters.3 Following Baxendale, they have proposed that the hydrogen peroxide-cerium(1V) reaction proceeds via the two-stage process Ce(1V)
+ HzOz
kl Ce(1II)
+ HO2 + Hf
(1)
k-1 k2
Ce(1V) 3- HOz +Ce(II1)
+ O2+ H+
(2) with Ic-llkz = 0.129 f 0.013 a t 0'. According to Baer and Stein, on the other hand, the reverse of reaction 1 does not occur to any significant extent.4 The conclusions of Baer and Stein are based on studies of the stoichiometry of the hydrogen peroxide-cerium(IV) system. The formation of perhydroxyl radicals as intermediates in the hydrogen peroxide-cerium(IV) reaction, as required by the above reaction scheme, recently has been confirmed by means of electron spin resonance spectroscopy.K16 (1) Research performed under the auspices of the U. S. Atomic Energy Commission. (2) J. H. Baxendale, J. Chem. SOC.(London), Spec. Publ. No. 1 , 40 (1954). (3) P. B. Sigler and B. J. Masters, J . A n . Chem. SOC.,79, 6353 (1957). (4) S. Baer and G. Stein, J . Chem. Soc., 3176 (1953). (5) E. Srtito and B. H. J. Bielski, J. Am. Chem. Soc., 83, 4467 (1961). (6) B. €1. J. Bielski and E. Balto, J. Phys. Chem., 66,2266 (1962).
202
Vol. 67
NOTES
10
7.8
7
c
9
8
.a Io
k
6
a
1
1
2
3 I .4
0
20
IO
Fig. 2.-Dependence
,
I .o L' 0
0
30
40
IO3 [H20,],/ [Ce (DI)], ,
I .2 1.1
4
0
2
4
6
I
I
8
IO
TIME)
12
I
I
i
1
14
16
18
20
1
M ILLISEC,
Fig. 1.-First-order plot of the kinetic data obtained in the P, first experiment of Table 11. [Ce(IV)lo = 92.5 X [H202]0= 7.50 X 10-0 P , [HzSOa] = 0.8 N , T = 25.0".
In order to obtain additional information, we have studied the kinetics of the hydrogen peroxide-cerium(IV) reaction in 0.8 N HzS04 using a flow technique. The results obtained confirm the mechanism proposed by Sigler and Masters, and in addition establish that kl = 1.0 0.1 X 106F-lsec.-lat 25.0'.
*
Experimental Ceric sulfate (G. Frederick Smith), cerous sulfate (Amend Drug tu Chemical Go.), hydrogen peroxide (Baker Analyzed Reagent), and sulfuric acid (Baker and Adamson) were used without further purification. The solutions for the kinetic measurements were prepared with triply-distilled water. The kinetics were studied in 0.8 N sulfuric acid; no attempt was made to keep the ionic strength of the solutions constant. The ceric sulfate solution was standardized spectrophotometrically; a value of 5580 was assumed for the extinction coefficient of ceric sulfate a t 320 mp.7 The cerous sulfate and hydrogen peroxide were estimated either directly or indirectly as ceric sulfate; the cerous sulfate was estimated after the addition of excess potassium persulfate, and the hydrogen peroxide after the addition of excess ceric sulfate. The disappearance of ceric sulfate was followed spectrophotometrically by means of a modified version of the rapid-mixing and flow apparatus which has been described previously.*P8 The modification allowed reactions with half-times down to a few milliseconds to be studied by means of the stopped-flow technique. All the kinetic measurements were made a t 25.0'. A first-order plot of the kinetic data obtained in the first experiment of Table I1 is shown in Fig. 1.
Results and Discussion If it is assumed that the hydrogen peroxide-cerium(IV) reaction proceeds via reactions 1 and 2, the rate of disappearance oi cerium(1V) is given by d[Ce(lV)l = {A-l[Ce(III)] - k2[Ce(IV)]}[HOnl dt kl [Ce(IV)1 [HzOZI (3)
where the hydrogen ion and sulfate ion concentrations have been included in the appropriate rate constants and the reverse of reaction 2 has been neglected. The (7)A. I. Medalia and B. J. Byrne, Anal. Chenz., 2 3 , 453 (1951). ( 8 ) N. Sutin and B. M. Gordon, J . Am. Chem. Soc., 83, 7 0 (1961). (9) G. Dulz and N. Sutin. t o be published.
of
kobsd
on [H~02],/[Ce(III)]~ at 25.0'.
TABLEI OXIDATION O F HYDROGEN PEROXIDE BY CERIC SULFATE [Ce(III)] >> [Ce(IV)], [HzOzl > [Ce(IV)], [HzSO~I= 0.8 N , T = 25.0" 10n[Ce(IV) lo, F
lO*[Ce(III)lo, F
lO.I[HzOs]o, F
15.5 33.0 33.0 35.5 35.5 35.5 35.5 35.5 35.5 36.5 36.5 50.0 64.0
6.20 6.40 2.40 5.93 5.93 5.93 5.93 5.93 5.93 5.93 5.93 6.20 6.40
4.08 5.04 5.04 7.65 10.3 12.7 15.5 18.2 20.4 2.57 5.02 4.08 5.04
lO-'kobad F-1 sec.'1
2.10 2.14 4.60 3.92 4.81 5.81 7.64 8.50 9.38 1.46 2.66 1.74 1.96
TABLEI1 OXIDATION OF HYDROGEN PEXOXIDE BY CERICSULFATE [Ce(IV)] > [HZOz],[H2SOa]= 0.8 N , T = 25.0' 106 [Ce(IV) 10,
1 0 8 [Ce(III) lo,
106 [HaO?lo,
F
F
F
92.5 103 147 147 147 147 147 142 147 147 147 147 200 200 200 200 294 294 294 294
.. ..
5.56 11.2 20.2 28.0 42.6 55.6 60.7 69.6 88.6 111.o 3.90 4.88 5.75 7.96 8.90 33.4 64.0 83.4
7.50 5.50 6.50 6.50 6.50 6.50 6.50 6.50 6.50 6 50 6.50 6.50 5.00 5.00 5.00 5.00 13.5 13.5 13.5 13.5
k'obad,
see. -I
101 94.5 24.7 15.1 9.30 7.58 5.33 4.74 4.00 3.49 2.88 2.25 88 9 71.9 62.5 53.6 71.0 28.0 17.1 11.6
steady-state equation for the concentration of the perhydroxyl radical is
h [Ce(IV) 1 [HOzl (4) Solving eq. 4 for the steady-state concentration of the pel-hydroxyl radical and introducing this value into eq. 3 gives
COMMUNICATION TO THE EDITOR
Jan., 1963 d[Ce(IV)] -- _ dt
-
2klh [Ce(IV) 1 [HzOz1 (5) Ic-l[Ce(III)] kz[Ce(IV)]
+
The rate of disappearance of cerium(1V) will be given by eq. 5 provided eq. 3 and the steady-state approximation for the concentration of the perhydroxyl radical are valid. It is possible to derive two limiting kinetic expressions, which depend on the relative concentrations of cerous sulfate, ceric sulfate, and hydrogen peroxide, and these expressions will be discussed in turn. Case 1.-In this case [Ce(III)] >> [Ce(IV)j and [H202]> [Ce(IV)]. If k-l[Ce(III)] >> kz[Ce(IV>] eq. 5 becomes d[Ce(IVIJ -
at
2kIlc2[H202][Ce(IV)Iz ?~-~[Ce(I11)]
- kobsd[Ce(IV)]'
I
I
I
I
I
7
6 ci c
5
LL v)
- 0 a 1
4
---.3 -w E
a
a
2
VI
0 I
0
[ce (mi], / [ c e
(6)
Fig. 3.-Dependence
where 2 k k z [%Os I k-l[Ce(lI1)]
203
(m)Im,
of [Ce(IV)]m/k'o,,saon [Ce(III)]~/[Ce(IV)lm a t 25.0'.
Values of k'obsd calculated from the integrated form of eq. 8 are presented in Table 11. As in case I, each valule of k',bsd is the mean of four determinations with the individual determinations differing from the mean by Values of kobsd calculated from the integrated form of less than 10%. The values of [Ce(IV)]/k',bpd are plotted eq. 6 are presented in Table I. Each value of kobsd against [Ce(III) lo/ [Ce(IV)Irn in Fig. 3, where [Cepresented in this table is the mean of four determina(IV)], is the mean concentration of cerium(1V) during tions. Individual determinations differed from the the run. It will be seen that the plot is linear and OF mean by less than 10%. The values of kobsd are plotted against [HzOzIm/[Ce(III)]o in Fig. 2, where [ H Z O ~ ] ~finite ~ ~ intercept, as required by eq. 10. Since the intercept is very small, the value of h1 was determined is the mean concentration of hydrogen peroxide during directly by measuriing the rate of the hydrogen per-. the run. It will be seen that the plot is linear, and of oxide-cerium(1V) reaction at high cerium(1V) concenzero intercept, as required by eq. 7. The value of trations, but in the absence of added cerium(II1) lcllc2/k-1 a t 25.0' calculated from the slope is 1.37 X The results of these measurements, which are included lo7F-l sec.-l. in Table 11, give kl = 1.0 i: 0.1 X 106F-l sec.-l at Case 11.-In this case [Ce(IV)] > [HzOz]. Under 25.0°. This value of kl is in good agreement with the hhese conditions eq. 5 reduces to value given by the intercept of Fig. 3. The slope of Fig. 3 gives 1.19 X lo7F-I sec.-l for the value of klkz/ k-1 at 25.0', which if; in good agreement with the value calculated from the slope of Fig. 2. The mean of where these values is 1.28 =t 0.1 X lo7F-l sec.-I, corresponding to a value of 13 f 2 for kB/k-l a t 25.0'. This value is consistent with that obtained by Sigler and Masters a t Oo.3 The studies described above are thus consistent with the mechanism proposed by Sigler and Masters and in addition establish that kl = 1.0 f 0.1 x lo6 F-l sec.-l in 0.8 N sulfuric acid at 25.0'. kobsd
=
(7)
COMMUNICATION TO THE EDITOR O N L4RGE DIFFERENCES I N T H E EQUILIBRIUM SOLUBILITIES OF HYDROGEN AND DEUTERlUM IN PLATINUMPALLADITJM ALLOYS
Rir: The absorption of deuterium by palladium has been investigated recently, utilizing the technique of direct absorption of deuterium gas from acidic heavy water solutions.l The Of absorption by (1) T. B. Flanagan, J Phyls. Chem., 06, 280 (1961).
changes of relative rcsistance and electrode potential of the wire specimens It previously has been demonstrated that equilibrium absorption data are obtained with this technique.2 I n agreement with earlier work, it was found that the final equilibrium solubility of deuterium in palladium does not differ appreciably from that of hydrogen in palladi~m,i.e., D/Pd = 0.65, H/Pd 0.69 (atomic,ratios, 250, 1 atm. pressure). Having regard to this difference in solubilities in E
( 2 ) T. €3. Flanagan and F. A. Leuls, Trans. Faladay Sac., 66, 1409 (1959).
