T H E CATALYTIC DECOMPOSITIOK OF FORMIC -1CID 1-hPOUR BY C. H. D. CLARK AND B. TOPLEY
The catalytic decomposition of formic acid vapour at the surfaces of metals, glasses, silica and metallic oxides can be explained on the view of contact catalysis, according to which the reaction takes place in an adsorbed layer one molecule thick, partially covering the catalyst surface. The independence of the two alternative modes of decomposition on glass, viz., ( I ) HCOOH+H2 COS,and ( 2 ) HCOOH-+H20 CO, when considered in connection with the large temperature coefficients of these two reactions, makes it very probable that the interpretation in terms of the ‘adsorption’ mechanism is correct in this instance.‘ The simplest assumption to explain the fact of catalysis by surfaces is that the energy necessary for activation of the reactant molecules is lowered by association with the catalytic surface, but this leads one to expect a rough parallelism between temperature coefficient of reaction rate and specific catalytic activity (Le., the catalytic activity per unit area of superficial surface). Such a parallelism might, of course, be completely masked by a wide variation in the number of molecules adsorbed per unit area; but a przorz it would seem reasonable to expect that when the heat of adsorption is large, a considerable alteration of the activation energy would be produced, because of the ‘molecular field’ in which the adsorbed molecule finds itself, and, simultaneously, a relatzvely large fraction of the surface would be covered by the adsorbed layer. Further, the positzve correction to be applied to the empirical ‘energy of activation’ in kinetically unimolecular reactions, on account of decreasing adsorption with rise in temperature, is equal to the heat of adsorption. If we write kau. e+x RT, where
+
+
k = rate of reaction per unit area of catalytic surface, under standard conditions of concentration of reactant vapour and volume of reaction bulb; X = empirical energy of activation obtained from the temperature coefficient of reaction rate; X = heat of adsorption per gram-molecule; R = gas constant; u = fraction of surface covered by a unimolecular adsorbed layer, at temperature T ; then we might expect, in passing from one catalyst to another, that the larger values of would be associated with larger values of X and smaller values of A, so that qualztatztely there should be a correlation between e - 4 RT and k. Actually, there appears to be no correlation between these two experimental quantities. ‘Hinshelwood and Topley: J. Chem. SOC., 123,
1022
(1923).
122
C . H. D . CLARK AND B. TOPLEY
Because of this, measurements of catalytic efficiency have been made for a number of metals, which have not been prcviously studied with the object of ascertaining whether their specific catalytic activities for one and the same reaction (HCOOH--1H2 COz) can bc related to any one property of the catalysts; for instance, on one view, it seemed possible that catalytic efficiency might be connected with the number of 'free' electrons in the metals, as measured by the excess of their specific heats over the value 3R. In the event, however, the results were not very helpful from this point of view, because in some cases the actual effects of the metals concerned seemed to be masked by an oxide film; but it has nevertheless seemed worth while to record the experimental results which have been so far obtained. The experimental method employed was similar to that used by Hinshelwood. The catalyst, together with a weighed amount of formic acid, contained in a narrow capillary with a very fine tip, was introduced into a bulb of known volume, which was cooled in ice, evacuated to below 0.1 mm. with a Hyvac pump, and sealed off. Vapour baths of anisol, aniline, nitrobenzene and quinoline were used to maintain suitable constant temperatures during the decomposition. The unimolecular velocity constant, reduced t o I sq. cm. of catalyst surface in a 2 0 C.C.reaction bulb, is taken as a measure of the specific catalytic activity. This constant will be denoted by k . Some experiments with osmium and tungsten were also carried out in a manometric apparatus exactly like the one already described.' Analyses of the gaseous products of decomposition were carried out in a Haldane apparatus. Catalysis by Osmium
+
The use of osmium as a catalyst was suggested by the observation of E. Muller and K. Sponsel? that, in respect of decomposition of aqueous solutions of formic acid, the activity of metallic rhodium decreases as the purity of the metal increases, and appears to be connected with traces of osmium in the rhodium. Metallic osmium was obtainable only in the form of a powder. .A rough estimate of the effective surface of this powder was made by determining the size of the grains. The particles were 'graded' by sifting through fine copper gauze, and the number of particles determined in two ways:( I ) By counting under a microscope the number of particles in a weighed amount of the powder; this gave 3.71 X ro6 particles per gram. (2) By the application of Stokes' Law, which gave I X ro6 particles per gram. From the mean of these results, the surface area of I gram of the powder is calculated to be 80.1 sq. cm. This represents the order of magnitude of the superficial area of the osmium catalyst. The experiments with osmium were done with 9 0 7 ~formic acid. For decomposition at temperatures in the neighbourhood of zoo°C, analysis of the 'Hinshelwood and Topley: J. Chem. Soc., 123, 1022 (1923:. Z. Elektrochemie, 28, 307 (1922).
2
CATALYTIC DECOMPOSITION O F FORMIC ACID VAPOUR
123
reaction products showed that only hydrogen and carbon dioxide were produced. The reaction follows the unimolecular law, as the results in Table I of a typical experiment in the manometric apparatus show.
TABLE I Pressure p (mm. of Hg)
Time t (minutes)
0.006jj 0.0064j 0.00679 o . 0067 I 0.00685
116 I97 267 314 349 440
20
40 60 80 IO0
End Point
The absolute value of the velocity constant was determined at 153.9OC, in a series of five sealed bulb experiments, the results being k = 0.0131, 0.0078, 0.0146, 0.00767 and o.oogo3. The average value of the unimolecular constant k o , 0.0104. The t,emperature coefficient was found by using the device of allowing the reaction t o take place partly at one temperature, and then to proceed to completion at a higher temperature, in the manometric apparatus. The results of experiment, are given in Table 11.
+
TABLE
11
Temperatures ("C)
kcoi
Ace? (cab.)
154.6 185.0
0,00477 0.0209
18,900
0.00241 0.0111
154.9 185. I
0.003I9 0.0133
18,400
The value of A may be taken as 19,000 calories per gram, to the nearest Using this, the absolute magnitude of the unimolecular constant per unit area for a 2 0 C.C. bulb, is calculated to be 9.2 x 10-2 a t ZOO'C. joo calories.
Catalysis by Nickel The nickel catalyst was used in the form of pure nickel foil. At the temperatures of the experiment (183' and 2 1 0 ~ )the formic acid vapour (99.6$&) reacts with the metal, hydrogen being formed; after the reaction, the bulb contained drops of a greenish liquid,-presumably a solution of nickel fornate. A t the same time, carbon dioxide is formed in the reaction bulb, probably as a result of direct catalytic decomposition at the metallic surface, and not by the decomposition of nickel formate, since there was no sign of a deposit of metal on the smooth surfaces of the foil.
C. H . D. CLARK AND B. TOPLEY
124
I n a typical experiment at position :
ZIO',
the gas in the reaction bulb had the com-
H = 64.35% C 0 2 = 33.7%
co*=
2.0%
The 2.0% of carbon monoxide corresponded to the activity of the walls of the reaction bulb, which was of soda glass. Similar results were obtained in other experiments, though the ratio of carbon dioxide to hydrogen is somewhat variable, and increases at the lower temperature. Under the conditions of these bulb experiments, there is no formation of methane by reduction of carbon dioxide by hydrogen, as observed by Sabatier with a finely divided nickel catalyst. Westcott and Engelder,' also obtained no methane when decomposing formic acid vapour by passing it over a nickel catalyst between zooo and joo'. On the assumption that the carbon dioxide is formed by the usual unimolecular reaction, independently of the simultaneous interaction of the nickel and formic acid vapour, the following rough values of kco, (calculated with natural logarithms, the unit of time being the minute) are a measure of the specific catalytic activity of nicke1:kco, per
I
3 X
183'C
I.6
2 IOOC
X
sq. em. 105 104
For comparison with other catalysts, the absolute magnitude of the unimolecular constant for nickel a t zoo°C may be taken as 1.0 X IO-^. Catalysis by Tungsten The catalyst was used in the form of a thin tungsten sheet. Catalytic decomposition takes place at a rate convenient for measurement a t ZIO'C;a manometric experiment proved that the reaction follows the unimolecular law. The gaseous reaction product was practically pure carbon monoxide. During the reaction, the bright surface of the tungsten becomes dull and slighily blue. Since all other metals that have been experimented with catalyse the carbon dioxide reaction mainly, it seems likely that the explanation of the production of carbon monoxide only in this case is that a thin film of the blue oxide of tungsten is rapidly formed on the surface of the metal, which then produces the usual catalysis by oxides yielding carbon monoxide. This would account for the discolouration of the metal, and, moreover, an experiment in which the yellow oxide of tungsten was heated at ZIOOCin a sealed bulb with formic acid vapour resulted in the reduction of the yellow to the blue oxide. -
J. Phys. Chem., 30,476 (1926).
CATALYTIC DECOMPOSITION O F FORMIC ACID VAPOUR
125
The existence of such a thin film of oxide on tungsten is probable, according to U. R. Evans,l in view of the marked passivity of the metal. The reaction bulb was evacuated to rather less than o . mm. ~ of residual air, which would leave sufficient oxygen to furnish an oxide film thick enough to produce interference colours. An interesting point is the unusual catalytic activity of this oxide film; the unimolecular constant kco per sq. cm. of superficial area of oxide layer is 2 . 5 X IO-^ at 21o’C. This is rather higher than the constant LO, for silver at the same temperature, and certainly very much greater than the specific activity of ordinary granular “blue” oxide of tungsten; possibly the association of the coherent oxide film with the underlying metal alters its properties. Catalysis by Molybdenum and Tantalum Both molybdenum and tantalym catalyse the decomposition, but the COS is not the main effect. A typical analysis usual reaction HCOOH+HZ of the gas obtained with molybdenum was:-
+
About one-third of the carbon monoxide and a little of the carbon dioxide in this experiment would have been due to the catalytic action of the soda glass of the reaction bulb; the excess of hydrogen is thus very marked. The interaction of formic acid and molybdenum is being further investigated. Catalysis by Silica, Pyrex and Soda Glass I n the course of the experiments on metals, an attempt was made to find a material for the reaction vessel which would have a negligibly small catalytic activity. The specific activity of silica and of Pyrex was compared with the products are mainly carbon monoxide and that of soda glass. At ZIOOC, water in each case. The means of the results of many experiments are represented by the following figures:kco
Soda Glass Pyrex Fused Silica
x 1.5 x 1.0 x 4
10-6 10-6
10-6
Table I11 contains the quantitative data available as to the catalytic decomposition of formic acid vapour on the surfaces of different catalysts; the values of k are calculated in reciprocal minutes, for unit surface of the catalyst in a reaction bulb of 20 c.c., using natural logarithms. “hletala and Metallic Compounds.”
120
C. H. D. CLARK AND B. TOPLEY
TABLE I11 Catalyst
kcoi (at zoo°C, per unit area).
Duroglass Gold SiIver Nickel Platinum Rhodium Osmium (Palladium
2.5
X
IO-'
9.2
x
10-6
I .03
I
4.4 2.4
9.2
10-3
X IO-^ X 10-~
x x X -
104
IO-^ IO& IO+)
kco
(at ZIO'C, per unit area).
Silica Pyrex Soda Glass Tungsten ( f oxide film)
x x 4 x 2.5 x I
10-6
1.5
10-6
10-6
IO-^
Our thanks are due to Dr. J. N. Williamson, of t.he Fuel Department, for the loan of a Haldane gas analysis apparatus. The University, Lee&, Sept. 24, 1967.
