J. CHARLES, P. W. KOPF, AND S. TOBY
1478
The Reaction of Pyrophoric Lead with Oxygen
by J. Charles, P. W. Kopf, and S. Toby' School of Chemistry, Rutgera, The State University, New Brunswick, New Jersey (Received October 220, 1966)
08903
Pyrophoric lead was prepared by the thermal decomposition of lead citrate in vacuo. The properties of the pyrophoric lead, as measured by stoichiometry of oxygen uptake, rate constant, and esr signal, were independent of preparation temperature above 400". Pyrophoricity was found in samples with preparation temperatures much above the melting point of lead, showing that the material cannot consist simply of finely divided lead. Analysis showed approximately 20 wt % of carbon was present, but it was found that the carbon does not take part in the oxidation. First-order disappearance of oxygen was found over an initial pressure range of 0.02 to 1 torr and a range of temperatures from -115 to 75'. The rate constant obeyed the Arrhenius law, k = (3.1 f 1) exp[(2900 f 200)IRTJ sec-I. The large negative entropy of activation is comparable to the over-all entropy of reaction.
Introduction Lead metal in a form which combusts spontaneously when exposed to air has been known since the last century. In the last 50 years three references2-* to the preparation and properties of pyrophoric lead have appeared. No serious attempts have been made to find the conditions which give the most active material, nor are there any quantitative data published on the chemical composition or reactivity of pyrophoric lead. The uptake of oxygen by pyrophoric lead is extremely rapid at room temperature, presumably because the metal is extremely finely divided. The kinetics of the reaction, a t least in its initial stages, give an insight into a metal oxidation process which is free of the usual complexities due to diffusion of oxygen through films of oxide. Since pyrophoric lead contains carbon, however, it is important to establish thatj the oxidation process involves the lead only and in this paper evidence is presented to show that the carbon plays no part.
complete when after several hours the pressure in the torr. system had fallen to less than Rates of oxygen uptake were measured by exposing pyrophoric lead to pure oxygen (Matheson Research Grade) and measuring pressure on a thermocouple gauge (A. F. Smith Co., Rochester, S. Y., RIodel 1013) which had been calibrated against a AIcLeod gauge. Rates of oxygen uptake were measured in a system of 580-cm3volume using approximately 0.25-g pyrophoric lead samples and a range of initial oxygen pressures from 0.02 to 2 torr. I n the initial experiments the pyrophoric lead samples were placed at the bottom of vertical tubes fitted with break-seals and arranged so that the oxygen would pass through about 25 cm of thermostated tubing before coming in contact with the lead. Occasional erratic results were obtained which were thought to be due to local heating accompanying the exothermic (52.5 kcal mole-') reaction. This was minimized by placing the pyrophoric lead in a thin layer in the horizontal portion of an L-shaped tube.
Experimental Section Lead citrate, formate, and tartrate were decomposed by heating approximately 50-g samples a t temperatures ranging from 300 to 600" in a vacuum system with continuous pumping. There was a large initial increase in pressure and the reaction was considered
(1) To whom correspondence should be addressed. (2) G.R. Levi and A. Celeri, Atti Accad. Naz. Limei, Mem., Classe Sei. Fis. Mat. Nat., IS] 7 , 350 (1928); Chem. Abstr., 22, 2299 (1928). (3) G. R. Levi and G . Rossi, Gazz. Chim. Ital., 68, 576 (1938); Chem. Abstr., 33, 1566 (1939). (4) N. I. Glistenko, T r . Voronezhsk. Gos. Univ., 42, 31 (1956); Chem. Abstr., 5 3 , 8903 (1959).
The Journal of Physical Chembtry
THEREACTION OF PYROPHORIC LEADWITH OXYGEN
Reproducibility improved, and the fact that rate constants for oxygen uptake were found to be approximately independent of initial oxygen pressure was taken as evidence that thermostating was satisfactory. This point is returned to in the Results. Carbon and lead determinations were carried out as follows. Air was slowly admitted to a weighed sample of pyrophoric lead. After reweighing, the sample was shaken with 1 : 1 nitric acid. The mixture was filtered and the washed residue, shown by simple chemical tests to be pure carbon, was weighed. The filtrate was analyzed for lead gravimetrically as sulfate. Electron spin resonance (esr) spectra were measured a t 350 Mc/s.
Results
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6 100 g.2
i
50
o
t
0--e--8
0 - 0 - 0-
A i
300
400 500 Preparation temperature, OC.
Figure 1. Reactivity of pyrophoric lead as a function of preparation temperature: A, per cent efficiency as a getter; B, first-order rate constant for uptake of oxygen; and C , electron spin resonance signal, spins per gram.
Lead formate, when decomposed at 325", gave massive lead and carbon. Lead tartrate decomposed over the range 300-425" gave a feebly pyrophoric material. Lead citrate gave an active pyrophoric material when decomposed in the range 300-600", and all subsequent experiments were performed with pyrophoric lead prepared from citrate. The product was a black powder. Analysis of several samples of pyrophoric lead gave the percentage by weight of lead as 83.5 f 1, the remainder being carbon. The percentage gain in weight on exposure to oxygen was 5.6 i 0.4. In Figure 1 some properties of pyrophoric lead are plotted against temperature of preparation. The efficiency of pyrophoric lead as an oxygen getter was calculated by assuming that all oxygen removed was due to Pb
+
'/202
+PbO
The actual weight of oxygen taken up divided by the stoichiometrically expected amount gave the getter efficiency, and this was plotted in Figure 1A. The rates of oxygen uptake were found to follow first-order kinetics. The rate constants for room temperature uptake are plotted as a function of preparation temperature in Figure 1B. The esr spectra of pyrophoric lead samples sealed in Pyrex tubes in vacuo were measured relative to a DPPH standard. The results are plotted as a function of preparation temperature in Figure 1C. In addition, oxygen was admitted to some samples; after oxidation had occurred, the excess oxygen was pumped off, and the esr spectra were reexamined. No significant difference was found. The validity of the first-order rate law was tested over a variety of initial pressures and amounts of oxygen removed. Figure 2A shows a first-order plot for
I
I
50
I
I
150 Time, seconds 100
I
200
Figure 2. Test of first-order rate law of reaction between pyrophoric lead and oxygen a t room temperature: A, varying initial pressures; and B, varying amounts of oxygen removed.
a 40-fold range of initial pressures. At the higher pressures a slight deviation from linearity occurred initially, but this disappeared after about 20 sec. The deviation may have been caused by a slight rise in temperature as the lead oxidized. As a precaution, initial oxygen pressures were kept below 0.25 torr when rate constants were measured at different temperatures. The constancy of the rate constant with oxygen uptake is shown in Figure 2B and shows little change Volume 70, Number 6
May 1066
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J. CHARLES, P. W. KOPF,AND S. TOBY
Table I : Distribution of Rate Constants from Different Samples of Pyrophoric Lead Range of k at 25' x lo2, sec-l No. of samples
0-0.5
0.6-1.0
1.1-1.5
1.6-2.0
2.1-2.5
2.6-3.0
3.1-3.5
3.6-4.0
4.1-4.5
0
2
5
5
8
5
1
1
1
up to about 4 X mole of oxygen removed. This corresponded to approximately 30% of the lead reacted. The reproducibility of rate constants measured from 28 different samples of pyrophoric lead is indicated in Table I. The values of k cluster around 0.0150.025 and samples giving rate constants in this range at 25' were used in studies a t other temperatures. The data shown in Figure 2B were obtained from two different samples and show that, for a given sample, good reproducibility was obtained. Using pyrophoric lead samples prepared at 450", rates of oxygen uptake were measured in the range -196 to 75', using a bath of the appropriate temperature around the sample. The oxygen uptake was fastest at -196', being so rapid that it was difficult to me:isure. We suspected that most of the rapid pressure decrease was due to adsorption. A sample of pyrophoric lead which had been exposed to air for several days was therefore exposed to oxygen at - 196", and the result is shown in Figure 3. This clearly represents adsorption, and the results a t - 196" were therefore discounted. Pressure-time curves a t - 140" similarly showed that most of the oxygen uptake was due to adsorption. However, at -115" and above, oxidation predominated over adsorption and could be corrected for. A pressure-time plot taken at -78" is shown in Figure 3. An adsorption plot is also shown and is of small importance, amounting to a correction of less than 5%. The adsorption that occurred was presumably due to the presence of carbon. As an additional check on possible interaction between the carbon and lead, a BET surface determination of pyrophoric lead before and after oxidation was made. The values obtained were 66 and 68 m2g-l, respectively, equal within the experimental error. The fact that the - 196" adsorption curve in Figure 3 is steeper than the reaction curves and does not follow first-order kinetics suggests that the reaction rate constants meavured from -115 to 75" have real significance. The rate constants obtained in the range -115 to 75" are shown as an Arrhenius plot in Figure 4. The parameters obtained are E = 2.9 f 0.2 kcal mole-' and A = 3.1. f 1 sec-l. The Journal of Physical C h a i a t r y
{
0.03
1
0
100 200 Time, sec.
300
Figure 3. Reactivity of pyrophoric lead prepared at 400': 0,first-order plots of oxygen uptake at -78"; and B, first-order plots of oxygen adsorption at -78 and 196' by ozidized pyrophoric lead.
-
Discussion Pyrophoric lead prepared by decomposition of lead citrate in vacuo a t temperatures above 350" is spectacularly reactive, giving sparks and flames when poured into air. The getter efficiency approaches 100% as indicated in Figure lA, showing that PbO was the only product of oxidation. This agrees with the observation of Levi and Celeri2 that X-ray analysis of pyrophoric lead exposed to air showed only P b and PbO. The esr intensities shown in Figure 1C are of interest. We obtained typical carbon resonance lines with a g value close to 2.00. Esr signals from carbon are well known,6 but we believe this is the first time that the esr signal from a pyrophoric metal has been noted. Ingram5* gives typical values of 5 X l O I 9 spins/g for an active charcoal prepared in vacuo, and this may be compared with our values of 1 X lo1*spins/g for pyrophoric lead containing 20 wt % carbon. Levi and (5). (a) D. J. E. Ingram, "Free Radicals as Studied by Electron Spin Resonance," Butterworth & Co. (Publishers) Ltd., London, 1968, pp 207-215; (b) I(. Antonowics, J . Chem. Phys., 36, 2046 (1962).
THEREACTIONOF PYROPHORIC LEADWITH OXYGEN
\
2.0
3.0 4.0 lOa/T"K.
5.0
6.0
7.0
Figure 4. Arrhenius plot of rate constant for reaction between pyrophoric lead and oxygen.
Rossi3 in an X-ray study found that pyrophoric lead had too large a particle size to show a pattern significantly different from that of massive lead. This contrasts with pyrophoric nickel, which is extremely finely divided. The importance of the carbon is emphasized by the fact that pyrophoric lead can be prepared at temperatures well above the melting point of lead without coalescing of the metal. Although pyrophoric lead contains about 20% carbon by weight, the percentage by volume may be as high as 90%. Thus, the lead appears to be strongly adsorbed on a carbon matrix, which prevents the lead particles from coalescing. However, the fact that the carbon plays no part in the oxidation process is indicated by the following: (a) the stoichiometry of oxygen uptake indicates that only PbO is formed; (b) the esr signal is typical for carbon and shows no change after the lead is oxidized; and (c) the BET specific surface is high (of the order of onetenth that of charcoal) and shows no change after oxidation. Relationship between Metallic Surface and Oxidation Law. The oxidation of a metal follows a rate law which depends on the state of the oxide formed and on the temperaturc6 At low temperatures most metals follow a logarithmic law of oxygen uptake with time. As the temperature increases, the state of the oxide and the rate of diffusion through the oxide may change. The oxidation law may then become cubic, parabolic, or linear. The "linear" law refers to a linear growth of oxide at constant oxygen pressure. This corresponds to the first-order law in oxygen found in the present
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work, for, if a given fraction of the collisions of the oxygen molecules with the lead result in reaction, then it follows that a t an oxygen pressure of p the rate a p a dp/dt so that a first-order law results. Even if the temperature is held constant, the oxidation law obeyed may depend on the state of subdivision of the metal. Copper a t 150" obeys a cubic law,6 but a study made on the low-pressure oxidation of copper foil which had been abraded in air and then reduced, gave a linear law7 in the range 140-300". A recent studya of the reaction of oxygen with evaporated films of lead gave a linear law at 230" and a complex logarithmic law in the range -78 to 100". The present work gives a linear law for finely divided lead in the latter range. It is possible that finely divided materials in general obey linear oxidation laws, for it is interesting to note that a recent investigation of the reaction between lead sulfide and oxygen gave a complex law for sulfide samples of low specific surface and a linear law for samples of high specific s ~ r f a c e . In ~ the case of the oxidation of metals, it is meaningless to make a comparison of rate constants for metals which obey different oxidation laws. I t is suggested that a study of the oxidation of metals in finely divided states will yield parameters which are directly comparable and which are functions of the metals rather than of the state of the oxide formed. Kinetic Parameters. The rate data obtained in the present study are for the early stages of the reaction, before lead depletion becomes important. This corresponds to a system containing an unlimited supply of available metal. The activation energy of 2.9 kcal mole-' found is considerably less than reported values for lead. I n the range 250-320" Weber and Baldwin's'O parabolic rate constants give E = 43 kcal mole-'. Anderson and Tare8 reported E = 11 kcal mole-' from the linear law obtained at two temperatures, 230 and 275". Our low value suggests that the restriction on reaction between pyrophoric lead and oxygen appears to lie in the entropy rather than the enthalpy of activation. If we assume the transition state theory formulation with a transmission coefficient of unity, then the observed A factor corresponds to A S = -55 cal deg-' mole-l a t 300°K. This large negative value is not un-
(6) 0. Kubaschewski and B. E. Hopkins, "Oxidation of Metals and Alloys," Academic Press Inc., New York, N. Y . , 1962. (7) F. J. Wilkins and E. K. Rideal, Proc. Roy. SOC.(London), A128, 394 (1930). (8) J. R. Anderson and V. B. Tare, J . Phys. Chen., 68, 1482 (1964). (9) L. J. Hillenbrand, J. Chem. Phys., 41, 3971 (1964). (10) E. Weber and W. M. Baldwin, J . Metals, 4, 854 (1952).
Volume 70,Number 6 May 1966
J. B. PERI
1482
reasonable since the transition state presumably resembles the products far more than the reactants.
o-----0
0 2
+
-Pb-Pb-Pb-Pb-
+-Pb-Pb.----.Pb-Pb- + 0
I
-Pb-Pb
0
I
Pb-Pb-
The entropy change for the removal of a mole of oxygen is ASo = -48 cal deg-’ mole-’, which is reasonably close to the value of AS* found. Finally, it may be mentioned that the rapid removal 30 sec) by pyrophoric lead sugof oxygen (half-life
-
gests its use for removal of oxygen from, for example, streams of other gases. Although pyrophoric lead does not react with oxygen as quickly as “flashed” barium metal,” it does react much faster than does sodium and is simpler to handle.
Acknowledgments. We wish to thank Rfr. E. Bretz and Dr. M. J. D. Low, both of this department, for help with the esr measurements and BET surface determinations. We are most grateful to the National Science Foundation for their support of this work through Grants GE-4048 and GE-7997. (11) See, for example, L. F. Ehrke and C. M.Slack, J . A p p l . Phys., 11, 129 (1940).
Infrared Study of the Reaction of Hydrogen Chloride with the Surface of y-Alumina and Its Effect on Surface “Acid” Sites*
by J. B. Peri Research and Development Department, American Oil Company, Whiting, Indiana
(Received October 2.5, 1066)
Chemisorption of HC1 on dry y-alumina and the effects thereof on surface sites of possible catalytic importance were studied by infrared and gravimetric techniques. Chemisorption affected preexisting hydroxyl (or deuteroxyl) groups selectively and formed new hydroxyl groups and water. Desorption of HC1 and water on heating normally left new hydroxyl groups plus chloride. All hydroxyl groups could, however, be removed by suitable treatment. The activity of HC1-treated alumina for isomerization of adsorbed l-butene increased with decreasing hydroxyl content, and the l-butene isomerized more rapidly than it exchanged hydrogen with residual deuteroxyl groups. ‘ l a sites” on alumina were affected by chloride so that carbon dioxide adsorbed thereon exhibited higher frequencies (2377 and 2392 cm-’ us. original 2370 cm-l). Changes in these sites probably explain changes in activity.
Introduction y-Alumina and other transition aluminas are cornmonly treated with halides to enhance their “acidity” and catalytic activity. The strong acid sites which exist on “dry” y-alumina appear to be ahminum ions The Journal of Phyeical Chemistry
(Lewis acids) partially exposed at a few special sites in the surface.2 A recent model explains how such (1) Presented at 150th National Meeting of the American Chemical Society, Atlantic City, N. J., Sept 1965.
