3842
T. M. GORRIE,P. W. KOPF,AND S. TOBY
The Kinetics of the Reaction of Some Pyrophoric Metals with Oxygen
by Thomas M. Gorrie,’ Peter W. Kopf,2and Sidney Toby School of Chemistry, Rutgers, The Skate University, New Brunswick, New Jer8ey 08903 (Received April 10, 1067)
Pyrophoric bismuth, cobalt, copper, iron, lead, nickel, and tin were prepared by the decomposition of the metal citrate, oxalate, or tartrate in vacuo at 350450O. The rates of oxidation were first order in oxygen pressure, and the rate constants gave linear Arrhenius plots in the range -120 to +75”. The entropies of activation were very similar to the standard entropies of the reactions with 1 mole of oxygen. A compensation effect was observed which was shown to be due to the fact that the free energy of activation was approximately constant at 18 kcal mole-’ for the oxidation of all of the metals studied.
Introduction The kinetics of the oxidation of metals is a subject of great theoretical and enormous practical interest. However, the rate-determining step in the oxidation of metals, as usually measured, is governed by diffusion processes in the solid phase. This complicates the kinetics considerably and the resulting rate laws have been described as linear, parabolic, cubic, or logarithmic, according to conditions.a A meaningful comparison of rate constants for metals which obey different rate laws during oxidation is not possible. In a previous paper‘ it was shown that the study of the reaction of pyrophoric lead with oxygen could give an insight into metal oxidation processes that was free of the usual complexities resulting from diffusion through oxide layers. In addition, an attempt was made to correlate the (kinetic) entropy of activation with the (thermodynamic) entropy of the reaction. This papers extends the investigation to other pyrophoric metals and describes what we believe to be the first comparison of the kinetic parameters for the oxidation of several metals.
Experimental Section The apparatus was similar to that previously described* except that a smaller reaction volume was used. In the course of repeating measurements on the oxidation of pyrophoric lead, we found an increase in the rate of oxygen uptake when the reaction volume was disconnected from associated tubing. This was presumably due to a viscous drag slowing the flow of oxygen. The Journal of Physical Chemistry
By eliminating dead space the reaction volume was reduced from its previous4 value of 580 cma to approximately 120 ,ma a t which the rate of oxygen uptake appeared to be independent of reaction volume. Tubing used throughout the apparatus had a minimum i.d. of 6mm. Approximately 0.5-g samples in break-seals were exposed to oxygen in the range -120 to 75”. Initial oxygen pressures were kept below 1.0 torr to minimize local heating and to conserve sample. Oxygen pressures were measured on a thermocouple gauge which had been calibrated against a McLeod gauge. It was noted that the calibration for oxygen was appreciably different from that for air. The most rapid pressure changes that could be measured with our apparatus corresponded to a rate constant of about 1.0 sec-I. This value was approached at the highest temperatures used. As previo~sly,~ rate constants were measured in the earlier stages of oxidation before the metal was exhausted. The pyrophoric metals were prepared by decomposing the appropriate salts at 350-450’ in a vacuum system with continuous pumping. All salts were commercially available and were used without further purification.
(1) NSF Undergraduate Research Participant, 1966-1967. (2) NSF Undergraduate Research Participant, 1965-1966. (3) 0. Kubaschewski and B. E. Hopkins, “Oxidation of Metals and Alloys,” Academic Press Inc., New York, N. Y., 1962. (4) J. Charles, P. W. Kopf, and 9. Toby, J . Phy8. Chem., 70, 1478 (1966).
REACTION OF SOME PYROPHORIC METALSWITH OXYGEN
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Table 1: Results of Decomposition of 20 Salts at 350450' in Vacuo Bi(II1)
Citrate
Cd(I1)
Ce(II1)
Pyro Carbon metal and massive metal
Co(I1)
Cu(I1)
Mmsive metal
Pyro metal
Formate
Mn(I1)
Fe(I1)
Green oxide and massive metal
Tartrate
Carbon and massive metal
Brown oxide
Pb(II1
Sn(I1)
Zn(II1
Pyro metal
Oxide
Pyro Pyro metal in metal poor yield Pyro Massive metal in metal poor yield
Red oxide
Oxalate
Ni(I1)
Pyro metal
Pyro Brownmetal red oxide
Feebly active metal
Pyro metal
Pyro metal
Table II : Kinetic and Thermodynamic Parameters for Oxidation of Metals AS *,
4/aBi 2co 2cu 4/3Fe 2Ni 2Pb 2Sn
Reaction
E, kcal/mole
Log A (A in PO-')
+ O2 + + + + 02 + 02 +
1 . 7 f 0.1 5.2 f 0.6 3.8 & 0.5 4.9 f 0.7 3.1 f 0 . 3 3.0 0 . 3 4.7 f 0.6
0.97 f 0 . 1 3.5 f O . 6 1.9 f 0 . 5 3.1 ztO.6 2.1 f 0 . 4 2.1 0.4 3.0 f0.7
2/aBi20a 2coo 0 2 -.+ 2 c u o 0 2 2/3Fez03 -.+ 2Ni0 4 2Pb0 0 2 2Sn0 -+
0 2
-+
Results Qualitative results of the decomposition of 20 different salts are given in Table I. Pyrophoric bismuth, cobalt, copper, iron, lead, nickel, and tin were prepared, and each metal showed the first-order uptake of oxygen expected from previous c~nsideration.~Although the reactivity was high in each case, pyrophoric copper and tin were found to lose their activity much more rapidly than the other metals. It is noteworthy that, as in the case for lead14pyrophoric bismuth and tin were formed a t temperatures considerably higher than the melting points of the metals. Arrhenius plots for rates of oxygen uptake are given in Figures 1 and 2. In all cases more than one sample of pyrophoric metal was used. The Arrhenius parameters with estimated errors are collected in Table 11. Also given are the standard entropy changes for the reactions per mole of oxygen. In the case of BizOs, Fe203,and PbO, the identity of the oxide was con-
oal/mols of
02 deg -56 It 1 -44 f 3 -52f 3 -46 f 3 -50 f 2 -51 It 2 -47 f 3
ASo, cal/mole of On dag
-43.0 -41.6 -44.1 -43.3 -45.0 -46.8 -46.6
firmed by gravimetric analysis with a correction for the carbon contained in the pyrophoric metal. For the other metals it was assumed that the stable oxide would be formed.
Discussion All of the pyrophoric metals described here have been made previously: but there are no accounts of quantitative kinetic comparisons. We were unable to make pyrophoric cadmium, cerium, manganese, or zinc by decomposition, but these metals have been made in pyrophoric form by hydrogen reduction methods. It may he significant that the enthalpies of reaction of the latter three metals with 1 mole of oxygen are considerably higher than those of the other metals. In the case of cadmium, appreciable sublimation of the ~
~~
compilation has been made by G . S. Bshn, Fall Meeting, Western States Section, The Combustion Institute, 1984, paper (6) A useful
64-31.
Volume 71, Number 18 Novernk 1967
T. M. GORRIE,P. W. KOPF,AND S. TOBY
3844
-13 I
-
c)
0
E-14 m
V Y
-15 m
+t
I
Q
I-
-16
-17 I
I
I
I
I
1
2
3
4
5
AH*, kcal mole-1 1
I
coI\
3
4
5
'(i I
6
7
Figure 3. Compensation effect: plot of TAS* us. AH for the oxidation of seven pyrophoric metals. The line drawn has a slope of unity.
*
1031T'K
Figure 1. Arrhenius plots for the oxidation of pyrophoric cobalt, nickel, lead, and tin.
0.1 -
-
.01-
23
v)
,001Y
t
I
Figure 2. Arrhenius plots for the oxidation of pyrophoric copper, bismuth, and iron.
metal was observed during the decomposition in vacuo, and the volatility of the metal probably precluded its formation in a finely divided state. The possibility exists that the observed rates were in part governed by gas-phase diffusion effects, and it was indeed found that the rate of oxygen uptake was The Journal of Physical Chemistry
measurably decreased by a constriction in the connecting tubing. This decrease affected the observed A factors but not the activation energies. If diffusion were paramount, one would not expect to find appreciably different activation energies for various metals. Although diffusion effects cannot be entirely ruled out, we do not think they play an important part in our data. The similarity between entropies of activation and entropies of reaction shown in Table I1 is quite striking. Although lessened somewhat, this similarity would still be apparent if appreciable quantities of a lower oxide were also formed [e.g., A.S"(Cuz0) = -35 cal (mole of O,)-l deg-'1. The restriction on reaction between metals and oxygen evidently is an entropy rather than an enthalpy effect, and the transition state resembles the products far more than the reactants. Our values of AS* are insufficiently precise for detailed speculation to be justified. However, there appears to be a slight increase in entropy in going froni transition state, which presumably consists of adsorbed oxygen molecules, to the oxide product. A comparison of the A factors and activation energies in Table I1 reveals that variations in the two parameters are parallel. For similar reactions it was noted by Hinshelwood and Legard6 that log A was proportional to E. This so-called compensation effect wm observed for solute diffusion in metals' and is to be (6) C. N. Hinshelwood and A. R. Legard, J. Chem. Sot., 587 (1935). R.H.S w a b , J . A p p l . Phys., 27, 554 (1956).
(7)
OXIDATION OF AQUOPENTAAMMINE
expected for a series of reactions involving similar bonds, Laidlers has considered the compensation effect for reactions in solution in terms of the free energy of activation, AG*. Taking this approach and assuming that AG* is approximately constant for the oxidation of all metals, then a plot of TA8* vs. AH* should be linear with a slope of unity. Using the simple transition-state theory formulation, we have evaluated TAS* (at, 298°K) and AH* from our Arrhenius pa-
3845
rameters, and the results are plotted in Figure 3. The data agree well with the line drawn with a slope of unity, suggesting that AG is constant for the oxidation of pyrophoric metals and possibly for all metals when a layer of oxide does not retard subsequent oxidation.
*
Acknowledgwnt. We are most grateful to the Nat,ional Science Foundation for support of this work. (8) K. J. Laidler, Trans. Faruduy SOC.,5 5 , 1725 (1959).
The Oxidation of Aquopentaammine and of Hexaamminecobalt(111) Ions
by Darwin D. Thusius and Henry Taube Department of Chemistry, Stanford Unitersity, Stanford, California 94306
(Receiaed April IO, 1967)
When &OB2- decomposes in the presence of C O ( N H ~ ) ~ O under H ~ ~ +the influence of Ag+ as catalyst, reduction of Co(II1) to Go2+ is observed with 02 as the major component of the gaseous product. Tracer experiments with O'*-enriched aquo ion show that in at least 0.7 of the acts leading to the reduction of Go2+, complex bound oxygen is transferred to the gas phase. Indirect arguments are presented which indicate that the external oxidizing agent attacking the Co(II1) complex is neither HO nor SO4-. Direct experiments using a solution containing Ag2+ show that this ion (or some species in labile equilibrium with it) oxidizes the Co(II1) coordinated water to oxygen. The water coordinated to Co(I1I) is shown to be more readily oxidizable than is other water in the system. The spontaneous decomposition of S2Og2- in the presence of either C O ( N H ~ ) ~ O H or~ Co~+ (NH3)83+leads to the production of Co2+. The rate of production of Co2f as the concentration of C O ( N H ~ ) ~ O H increases ~ ~ + reaches a limiting value which is very close to onehalf the limiting rate of production of HO by the spontaneous decomposition of S2082-. The dominant intrinsic gaseous product of the reaction is N2, a t least when Co(II1) is within 25% of the saturation value. Striking effects of substituting D for H are observed. These experiments taken together with others suggest that although &El3 is the coordination sphere of C O ( N H ~ ) ~ Ois Hoxidized, ~~+ HO does not attack the NH3 directly. The reaction is interpreted as involving oxidation of Co(II1) to the Co(1V) state, Co(IV) then being reduced by coordinated NH, by a 2e- change.
Introduction The effect of a metal ion on the reactivity of an associated ligand is becomhlg of increasing interest. A specialized part of this subject area is comprised of the reactions in which a ligand upon being oxidized
by an external oxidizing agent brings about reduction of the metal ion center to which it is attached.' One direction in which this interest is being pursued2 is (1) (a) P. Saffir and H. Taube, J . Am. Chem. SOC.,8 2 , 13 (19aO);(b) J. p. Candyin and J. Hallpem, ibid., 85, 2518 (1963).
Volume 71, Number 12 November 1967
