Activated Carbon as a Catalyst in Certain ... - ACS Publications

70

ELMER C. LARSEN AND JAMES H. W?LLTON

rule; sols containing relatively large amounts of stabilizing electrolyte did not.

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REFERENCES ( 1 ) BURTON AND BISHOP: J. Phys. Chem. 24, 701 (1920). (2) CUY,E. J.: J . Phys. Chem. 26,415 (1921). J. Phys. Chem. 39, 283 (1935). (3) FISHERAND SORUW: (1) FREUNDLICH AND SATHANSON: Kolloid-Z. 28, 258 (1921). (5) JUDDAND SORUM: J. Am. Chem. Soc. 62,2598 (1930). J. Am. Chem. See. 60, 1263 (1928). (6) SORUM: (7) WEISER:Hydrous O d e s . The McCraw-Hill Rook Company, Inc., Piew York

(1926).

ACTIVATED CARBON AS A CATALYST I N CERTAIX OXIDATION-REDUCTION REACTIONS' ELMER C. LARSEN

AND

JAMES H . WALTON

Department o j Chemistru, Universzty o/ Wisconsin, Madison, Wisconsin Received June 90, 1959

I. THEDECOMPOSITION OF HYDROGEN PEROXIDE Lemoine (17) in 1907 first reported that carbon is an effective catalyst for the decomposition of hydrogen peroxide solutions. I n 1923 Firth and Watson (5)found that heating ordinary sugar carbon in uucuo increases its activity greatly, but that after the reaction has proceeded for about 10 hr. the catalyst becomes ineffective. They also found (6) that the rate of decomposition is proportional to the quantity of catalyst used. It has since been observed (15) that heating in. an atmosphere of moist oxygen results in a still more active carbon. King (11) obtained a carbon of maximum activity on activating sugar carbon at OOOOC. and a product of minimum activity a t 45OoC. Hc found thc catalyst to bccomc incffcctive after the reaction had continued for about 90 min. Hc attributed thc dccay in activity to a chemical reaction bctwwn the pcroxidc and the surface oxide of the carbon. It is the purpose of this investigation to develop a technique for obtaining samples of activated carbon whose catalytic properties can be reproduced, and to use this carbon to obtain more accurate data on the factors affecting the rate of decomposition of hydrogen peroxide. Special emphasis is given to the decay of catalytic activity. In part I1 of this paper 1 This investigation was financed by a grant from the Research Committee of thr University of Wisconsin, Dean E . B. Fred, Chairman.

ACTIVATED CARBON A8 A CATALYST

71

the activity of carbon when used as a catalyst in two autoxidation reactions is discussed.


EXPERIMENTAL PROCEDURE

Preparation of carbon Mallinckrodt’s “reagent quality” sucrose was used in all the work reported here unless otherwise specified. Fifty-gram portions of sucrose were heated in silica dishes in a muffle furnace a t 5OO0C., and the charred residue ground to less than 100 mesh in an agate mortar. Ten grams of the ground carbon were placed in a long-necked quartz flask and heated in a muffle furnace a t 1000°C. until no more gases were evolved. The resulting carbon contained less than 0.04 per cent ash. The carbon was activated in an all-quartz system consisting of a large opaque silica cylinder into which was inserted a long-necked quartz tube with an activating chamber 15 cm. long and 4 cm. in diameter. Ten-gram portions of carbon were placed in the activating chamber, the system flushed out first with nitrogen, then with oxygen, and moist oxygen passed over the carbon at the rate of 5 ml. per minute. The oxygen train was the same as that used by King and Lawson (13)and consisted of one wash flask flled with water and two with a saturated solution of calcium chloride. The activation was continued for 12 hr., since King (11) found that for carbon activated at 800°C. the activity of the carbon becomes nearly constant after 10 to 12 hr. of heating. The temperature of the furnace was held constant to &5”C. At the end of 12 hr. the carbon was quickly poured out into small beakers, allowed to cool a moment in air, and then placed under vacuum until used. The loss in weight during activation represented about 40 per cent of the total carbon.

M&rials Merck’s “reagent quality” superoxol, which is free from stabilizer, was simply diluted to the desired concentrations with double-distilled water. These results agreed with those obtained using redistilled C.P. quality superoxol. All other materials used were of the highest quality obtainable, and were used without further purification. Apparatus The rate of decomposition of hydrogen peroxide was followed by measuring the rate of oxygen evolution in an apparatus similar to that described by Filson and Walton (4). Fifteen milliliters of peroxide solution were placed in the flask, and the carbon was suspended in a glass vial which broke immediately upon starting the run. It was necessary to apply a correction for the gases evolved on wetting the carbon. Unless otherwise stated, all work ww carried out at 25”C., using as

72

ELMER C. LARSEN AND JAMES H. WALTON


catalyst carbon activated a t 850°C. The results obtained in duplicate runs in experiments involjing carbon and peroxide only (runs in’which no foreign substances were added) agreed to within 2 per cent when the same batch of active carbon w&s used, and to within 5 per cent when carbon activated in two different batches was used. Concentrations are expressed as milliliters of available oxygen per 15 ml. of solution, corrected to standard conditions. FACTORS AFFECTING THE RATE OF DECOMPOSITION

Temperature of activation Carbon was activated a t approximately 75°C. intervals over the range from 335’ to 1045OC. Half-gram portions were allowed to react with 15 ml. of peroxide solution containing 75 ml. of available oxygen. The

volume of oxygcn evolved in the interval from 5 to 35 min. was taken as an indexcof its activity. The results are shown in figure 1. The carbon is inactive when activated below 35OOC.; the activity reaches a maximum for carbons activated a t 825-875OC. King (11) reported a pronounced minimum in activity a t 40OoC. and a large increase in activity a t lower temperatures. The activity of his “unactivated” carbon was equivalent

73


ACTIVATED CARBON .45 .4 CATALYST

to that of his carbon activated a t about 650' to 700°C. The present authors found that carbon which had not been activated possessed slight catalytic activity, and attribute this to activation which occurred when the sugar was charred. It is probable that the carbon used by King was prepared under such conditions that appreciable activation took place, and the activity he reports for low-temperature carbons is probably due to failure to reestablish the surface characteristic of lowtemperature carbon after the activation which occurred on charring. Figure 1 also gives the data of King (12) for the effect of tcmperaturc of activation on the cataphoretic velocity of activated carbon. The form of the two curves is practically identical. It is probable that the samc surface condition which gives rise to the high negative charge on the carbon is also that which is responsible for its catalytic activity in the decomposition of hydrogen peroxide. TABLE 1 Oxygen liberated from peroxide solution bu carbon f r o m different sources YOLUME OF OXTOEN

EVOLVED IN 5 TO 35

SOURCE OF CARBON

MINUTES

ml.

Gelatin ......... . . . . . . . . . . . . Immeasurably fast Lactose.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15.35 Dextrose. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14.35 ............................................... 11.40 Starch. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9.75 Acetylene black.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 0.68 _ I _ _ _ _

74



Initial peroxide concentration To determine the effect of the concentration of peroxide on the rate of decomposition, 0.3 g. of active carbon was allowed to react with 15 ml. of peroxide solution whose initial concentration was vaned from 15 ml. to 120 ml. of available oxygen. From figure 2 it is seen that the times of half-life and of third-life are proportional to the initial peroxide concentration. This is not a zero-order reaction, however, for the rate of the reaction decreases as the reaction proceeds. This decay in activity is discussed at length later in the paper. Since the amount of peroxide decomposed per unit time is proportional to the quantity of catalyst present (6),it is probable that the reaction takes place with peroxide adsorbed on the surface. If this is true, then the amount of oxygen evolved per unit time at different concentrations should be an index of the amount of peroxide adsorbed. It is then possible to calculate indirectly the constant 7t of Freundlich’s adsorption equation x/m = acn

Plotting the volume of oxygen evolved in 10 min. against the mean peroxide concentration gives a curve resembling an adsorption isotherm. Plotting the logarithm of the volume evolved against the logarithm of the mean concentration gives a straight line of slope 0.39. If 30 min. is chosen as the unit of time, the value of the slope is smaller because of the decay factor. If a smaller unit could conveniently be taken, a value greater than 0.39 would result. In later calculations the value 0.45has been used. If one attempts to determine reaction rate constants in the usual manner it is found that the reaction follows no apparent order of reaction until from 15 to 20 per cent of the peroxide has been decomposed. From 20 to 45 per cent decomposition the reaction appears to be a second-order reaction, and beyond 45 per cent decomposition it appears to be a fintorder reaction. This regularity was observed in all the reactions, regardless of the concentration of the peroxide, the amount of catalyst, or temperature of the reaction, providing that a fresh sample of catalyst was employed in each run. Variation of the quantity both of catalyst and of peroxide The quantity of catalyst and the initial concentration of peroxide were increased progressively from 0.075 g. of carbon and 15 ml. of available oxygen to 0.6 g. of carbon and 120 ml. of available oxygen. The ratio of carbon to peroxide was kept constant a t 0.1g. carbon to 20 ml. of available oxygen. The period of half-life is a linear function of the reciprocal of the initial peroxide concentration. This is shown in figure 3.

75

ACTIVATED CARBON A 8 A CATALYBT


Temperature of reaction (temperature coernent) In determining the influence of the temperature on the rate of peroxide decomposition, 0.3 g. of carbon WBS allowed to react with 15ml. of peroxide solution containing 50 ml. of available oxygen a t intervals of 5°C. from 15" to 35°C. The temperature coefficient, calculated from the reciprocal of period of half-life, WBS 1.90 over the range from 15" to 25"C., 1.97 from 20" to 3OoC., and 2.03 from 25" to 35°C. This compares with the value of 2.1 reported by Allmand and Style (1) for the thermal decomposition

I/C x

FIG.3

rd

PH

FIG.4

FIG.3. Variation of time of half-life with reciprocal of initial peroxide concentration when amounts of carbon and of hydrogen peroxide are increased in the same proportion. FIG.4. Effect of pH on time of half-lire. 8 represents runs using hydrochloric acid in place of sulfuric acid to adjust the pH in the acid range.

of hydrogen peroxide. The high value of the coefficient indicates that the rate of chemical reaction rather than the rate of diffusion is the limiting factor in determining the speed of reaction.

Hydrogen-ion concentrations The period of half-life for the decomposition of hydrogen peroxide solution containing 50 ml. of available oxygen and catalyzed by 0.3 g. of active carbon was determined over the pH range 2.6 to 9.5. Sodium hydroxide was used to adjust, the pH over the alkaline range, and sulfuric acid over the acid range, though it was found that hydrochlorio acid gave identical

76

ELMER

c.

LARSEN AND JAMES

n.

WALTON


results. The pH measurements were made with a glass electrode. The results are shown graphically in figure 4.

Metallic salts Salts of many heavy metals are in themselves effective catalysts for the decomposition of hydrogen peroxide. Iron salts (2) are known to be the most effective,but copper salts (14)are able to effect appreciable decomposition. In this work the catalytic activity ,of ferric sulfate and copper sulfate when used in conjunction with active carbon was determined. Samples (0.3 g.) of carbon and varying amounts of the salt were allowed

MILLIMOLES PER L. FERRIC SULFATE

MILLIMOLES PER L. COPPER SULFATE

FIG.5 FIQ.6 FIG. 5. Combined effect of ferric sulfate and active carbon on the rate of peroxide decomposition. FIG. 6. Combined effect of copper sulfate and active carbon on the rate of peroxide decomposition.

to react with a peroxide solution containing 50 ml. of available oxygen. The results are shown in figure 5 for ferric sulfate and figure 6 for copper sulfate. In both cases, the presence of small quantities of the salt reduced the total activity. This is probably due to the preferential adsorption of the salt by the carbon. As the concentration of the salt is further increased, the rate of decomposition begins to increase. With ferric sulfate this increase is very great and indicates a promoter reaction, since when ferric sulfate is used alone, the rate of increase of the speed of reaction tends to fall off over the same range of concentration.

ACTIVATED CARBON AB A CATALYBT

77

THE DECAY IN CATALYTIC ACTIVITY OF CARBON


The eflect of m o i s t u r eand oxygen The time of half-life w&s determined for the decomposition of 15 ml. of hydrogen peroxide solution containing 50 ml. of available oxygen, using as catalyst 0.3 g. of carbon which had been treated in a variety of ways. With no special treatment the period of half-life was 100 min. When the carbon had been moistened with water and dried in an oven for 6 hr. a t llO"C., the time of half-life w&s 108 min. A sample kept in a bomb a t an oxygen pressure of 40 atmospheres for 3 days gave a half-life of 114 min. When the carbon was moistened with water and placed in the bomb a t an oxygen pressure of 40 atmospheres, the time of half-life was increased to 148 min. Another sample was suspended in water and tank oxygen bubbled through the suspension for 8 days. The time of half-life was 173 min. The carbon is most readily deactivated by the combined action of moisture and oxygen. Maintaining constant concentration of peroxide To determine the rate of decomposition of hydrogen peroxide under conditions such that the peroxide concentration remained sensibly constant, 0.1 g. of carbon was allowed to react with 15 ml. of peroxide solution containing 817 ml. of available oxygen. During the course of the observations a total of 42 ml. of oxygen was evolved,-about 5 per cent of the total available quantity. The rate of decomposition over the same period decreased from 2.09 to 0.22 ml. of oxygen per minute. Compared to this enormous decrease in activity, the 5 per cent decrease in peroxide concentration is almost negligible, and can be regarded as remaining constant throughout the experiment. In figure 7 the variation of both the total amount of oxygen evolved and the rate of decomposition with time are given. In determining this rate of decomposition, a correction factor of 0.01 ml. per minute was applied to allow for the spontaneous decomposition of the concentrated peroxide solution. At the end of the run, 422 ml. of oxygen had been decomposed per gram of carbon, and decomposition was still taking place at the rate of 2.2 ml. of oxygen per gram of carbon per minute. The activity was decreasing very slowly over this range. King (11) has postulated that active carbon is not a true catalyst, but that a reaction occurs between active carbon and the high-temperature surface oxide, converting it into the low-temperature type with the liberation of oxygen. However, he reported complete decay in activity of 850°C. carbon when only 70 ml. of oxygen had been evolved per gram of carbon. I t is possible that the surface oxide is changing during the reaction, but in light of the

78



data presented here, the simple stoichiometric change postulated by King is not sufficient to explain the observed facts.

Repeated use of the same carbon The rate of decomposition of a hydrogen peroxide solution containing 50 ml. of available oxygen with 0.3 g. of carbon as catalyst was followed in the usual manner. Then this same carbon, whose activity had greatly decreased, was allowed to react a second and then a third time with the same quantity of hydrogen peroxide. In another series the procedure was

TIME IN MINUTES

TIME IN MINUTES

FIG.8 fiQ.7 FIQ. 7. Rate of decomposition with large exceea of hydrogen peroxide; 0 , total volume of oxygen; c ) , dc/dt. FIQ. 8. Log c ue7.3~8time for three consecutive reactions using the same catalyst

changed. The carbon was allowed to react with a solution containing 100 ml. of available oxygen. When this carbon was used a second time, its activity was comparable with that of carbon used in the third run described above. The total quantity of hydrogen peroxide decomposed is then the important factor. In figure 8 the logarithm of concentration is plotted against time for these three consecutive reactions. The first reaction was carried to 80 per cent completion (4hr.), the second to 68 per cent completion (9 hr.), and the third to 45 per cent completion (9 hr.). I n each case they appear to follow a first-order reaction during the greater part of the decomposition. There waa no tendency to deviate from the straight line for as long a time


ACTIVATED CARBON AS A CATALYST

79

as the course of the reaction wm followed. However, although these linear relationships are found to hold within a run, the values lor the reaction rate constants for the three runs differ widely. For the reaction in which fresh carbon was used the first-order reaction constant was 97 X 10W, for the second run 40 X and for carbon used a third time 18 X reciprocal seconds. In the calculation of the reaction rate constants just given, no consideration was given to the influence of adsorption. When a correction is made

.475

.oeo

L

,045

,034

,415

TEMPERATURE t FIG.9 FIG.10 FIG.9. Variation of first-order reaction constant, corrected for adsorption, with amount of hydrogen peroxide decomposed in three consecutive reactions. FIG.10. Effect of temperature of activation of carbon on adsorptive capacity for iodine. 0 , 795°C. carbon left in water 12 hr. and dried; @, same left in hydrogen peroxide solution.

for adsorption, the equation for the first-order reaction constant becomes (18) ch-" - e:-" = kt

where 12 is the Freundlich adsorption constant. Plotting C L - ~ against, time should give a straight line of slope k. Calculations were made for the series of three consecutive reactions, using for n the value 0.45. For these calculations the concentration of peroxide w&s expressed in moles per liter, and time in seconds. These calculations give a continuously decreasing value for the first-order "constant". In figure 9 the values of the constants thus calculated are plotted

80

ELMER C. LARSEN AND JAMES R. WALTON

against the total volume of oxygen evolved during the reactions. Though in each of the three trials there is a sharp decrease in the value of the constant during the first'part of the reaction, taken as a whole the values so obtained show a continuous decay in catalytic activity rather than the three sharp breaks obtained when calculations are made without taking adsorption into consideration. Downloaded by CENTRAL MICHIGAN UNIV on September 10, 2015 | http://pubs.acs.org Publication Date: January 1, 1940 | doi: 10.1021/j150397a009

ADSORPTION OF IODINE FROM AQUEOUS SOLUTIONS

Iodine has been used extensively to determine the relative surface area of carbon because it has been assumed (11) that not only are the molecules small enough to enter the ultra-fine pore structure, but that, unlike acids and bases, it is adsorbed equally well by high- and low-temperature carbons. A determination of the adsorption capacity for iodine of carbons activated a t different temperatures seemed advisable. The adsorption experiments were carried out according to the method of Kruyt and deKadt (16), using 0.5 g. samples of carbon and 50 ml. of 0.1084 N iodine solution. The results, calculated on the basis of grams of iodine adsorbed per gram of carbon, are shown in figure 10. There is a minimum in the amount of iodine adsorbed by carbon activated a t 500°C. This minimum occurs a t the same temperature of activation for which Schilow (20) reports a minimum in acid adsorption. It is scarcely probable that increasing the temperature of activation would decrease the total surface area. It would indicate, rather, that the nature of the carbon surface influences the amount of iodine adsorbed as well as the amount of acids and bases. To one sample of 795OC. carbon, distilled water was added; to another some stabilizer-free hydrogen peroxide. Both samples were allowed to stand overnight, then filtered and dried in vacuo a t 65OC. for 6 hr. The carbon that had been kept in water adsorbed the same quantity of iodine as before this treatment, while that which had reacted with the peroxide adsorbed considerably less. The total surface area in each case should be the same. King (11) has postulated that the reaction with peroxide converts the surface oxide of the carbon from the high-temperature to the low-temperature form. This decrease in the amount of iodine adsorbed after a reaction between peroxide and carbon lends further support to the postulate, and indicates that the nature of the surface influences the amount of iodine adsorbed. One would scarcely expect the amount of adsorption to be equivalent to that of low-temperature carbon, for activation a t high temperatures would increase the surface area. The large increase in iodine adsorption for high-temperature carbons is caused by the combined effect of increased surface area and the formation of a new surface oxide.

ACTIVATED CARBON A 8 A CATALYBT

81


DISCU88ION OF RESULTB

In spite of the regularity observed in the rate of decomposition of hydrogen peroxide when catalyzed by active carbon, the reaction is too complex to permit a simple quantitative analysis of the observed facts. In addition to the complications due to considerations of adsorption, there is superimposed an additional complicating factor,-the decay in activity of the carbon as the reaction proceeds. Most investigators now favor the hypothesis that the specific properties of active carbons are a result of the surface oxides formed a t different temperatures of activation (19). At least two types of surface oxides are believed to exist. Schilow (20) postulates three types of surface oxides,A, B, and C. At temperatures above 8OO0C., oxide B alone is present. At lower temperatures the oxide B is replaced by C, but B is present again in increasing amounts below 500°C. Since 85OOC. carbon is most effective in the decomposition of peroxide, oxide B must be the catalyst. If this is true, then the activity should increase for carbons activated below 5OO0C., but this does not occur. The existence of only two oxides as postulated by King,-an alkaline oxide formed at high temperatures and an acid oxide formed at low temperatures,---explains the observed facts much better. The data of King on the pH of charcoal suspensions and the measurements of iodine adsorption reported here support the view that deactivation of the carbon and a change in the nature of the surface occur simultaneously. However, contrary to the reports of previous investigators, complete destruction of catalytic activity was not observed, even after the reaction had continued for more than 20 hr. The carbon used in the extremely concentrated peroxide solution was still active even after it had decomposed six times the amount of peroxide that completely deactivated the carbon used by King. This indicates that the stoichiometric change in surface oxide does not adequately describe the change that occurs. Qualitatively, at least, these results can be rationalized by picturing the reaction as proceeding through a chain mechanism, the length of chain increasing with increasing peroxide concentration. Haber and Weiss and coworkers offer a variety of evidence to show that the catalytic decomposition of hydrogen peroxide by enzymes (9, 23), by various metals (23), and by ferrous and ferric ions (8) proceeds through a short chain mechanism involving the radicals OH and HOs. If the chain. mechanism is assumed, then as long as the concentration of the peroxide remains constant, the rate of decomposition should be proportional to the quantity of catalyst present. Increasing the concentration increases the length of the chains and more peroxide molecules are decomposed for a given number of initial chains. Thus, even if the initia-

82

ELMER C. LARSEN .4ND JAMES H. WALTON


tion of a chain results in a portion of the surface being changed from the alkaline to acid type of oxide, the actual quantity of peroxide decomposed would depend upon the concentration of the peroxide and would not be fixed, &s King suggests. The effect of concentration on the length of chain would also explain the increase in velocity constant observed each time a fresh portion of peroxide was added to a partially deactivated sample of carbon.

11. THEAUTOXIDATION OF STANNOUS CHLORIDE AND POTASSIUM URATE In part I of this paper the activity of activated carbon was measured by its effect on the rate of decomposition of hydrogen peroxide. In this part of the paper its activity will be measured by its effect on the rate of autoxidation of stannous chloride in an acid solution and of potassium urate in an alkaline solution. Apparatus and procedure The rate of autoxidation of these substances was followed by observing the decrease in volume of oxygen above the solution, using the method described by Filson and Walton (4). The stock solution of stannous chloride (Merck’s “reagent quality”) was 0.354 N with respect to stannous chloride and 0.8 N with respect to hydrochloric acid. The stock solution of potassium urate contained 20 g. of uric acid (Merck’s) per liter and 25 g. of potassium hydroxide per liter. Both solutions were kept under an atmosphere of nitrogen. Samples (0.5 g.) of carbon were allowed to react with 15 ml. of the stock solution a t 25OC. For the stannous chloride the rate of autoxidation in air wm determined; for potassium urate the rate of autoxidation in oxygen. Autoxidation of stannous chloride Haring and Walton (10) found that ordinary powdered willow charcoal accelerates the rate of autoxidation of stannous chloride in acid solution very greatly, but no work was done on the effect of activated sugar carbon. In this investigation carbons activated a t temperatures ranging from 350’ to 85OOC. were employed as catalysts. In figure 11 the speed of reaction, expressed as the reciprocal of the time required for oxidation of half the stannous chloride present, is plotted against the temperature of activation of the catalyst. It is seen that the speed of reaction is greatest when carbons activated in the neighborhood of 575°C. are used as catalysts. The rate of reaction decreases rapidly when carbons activated a t temperatures above 650°C. and below 450°C. are used. I t was found that this reaction is a zero-order reaction. There was no indication that the activity of the catalyst decays as the reaction proceeds, as is the case in the decomposition of hydrogen peroxide.


ACTIVATGD CARBON .iS A CAThLYST

53

I t is interesting to compare these results with those obtained by Schilow (20) and coworkers for the adsorption of acids and alkali on carbons activated a t different temperatures. The results they obtained for the adsorption of hydrochloric acid are also shown in figure 11. They found a maximum adsorption of alkali and a minimum adsorption of acids on carbon activated a t 540°C. Thus it is seen that those carbons which are capable of adsorbing the least amount of acid are those which are most effective in catalyzing the autoxidation of stannous chloride. The decrease in rate of autoxidation when carbons activated a t higher or lower tempera-

TEMPERATURE "C. FIQ. 11. Effect of temperature of activation: 0, on the rate of autoxidation of stannous chloride (work of the present authors); (3, on adsorption of hydrogen chloride (from Schilow).

tures are used can be attributed to the preferential adsorption of acid from the acid solution of stannous chloride, the acid occupying surface which would otherwise be available for stannous chloride. Autoxidation of potassium urate Frbrejacque (7) found that, using commercial activated charcoal in the ratio of 1 part of charcoal to 20 parts of water, the time required for the autoxidation of a quantity of potassium urate was reduced from 15 hr. to 10 min. No work was reported for the effect of activated sugar carbon. In this work carbons activated a t temperatures from 350' to 850°C. were


84

kLMEH C. LARSEN AND JAMES

H. WALTON

employed as catalysts. In table 2 the time required to oxidize half the potassium urate present in solution is recorded for each type of carbon. These experiments were difficult to reproduce. They were carried out in triplicate and the mean of the three trials recorded. The individual runs differed from each other by as much as 10 per cent. The reaction was a zero-order reaction, and was independent of the rate of shaking the solution. As in the autoxidation of stannous chloride, here too there was no indication of decay in catalytic activity of the carbon. Approximately 2 moles of oxygen were absorbed for every 3 moles of potassium urate. This agrees with the observations of Fr&ejacque, who found that between one and two atoms of oxygen are needed to oxidize one molecule of uric acid. TABLE 2 Effect of temperature of activation on the rate of autoxidation of potassium urate T I Y P E U T U R E OF ACFIYATION

I

*C.

TlYE FOR HALF-OXIDATION

minu&

No carbon Unactivsted 360

133

67 32.5 30

430

n

500 570

24.5 28 21.5 19.5 20

650 730 800

860

From table 2 it is seen that the rate of reaction increases as the temperature of activation is increased. There is no range of pronounced activity, as was observed for both the decomposition of hydrogen peroxide and the autoxidation of stannous chloride. Truszkowski (21) found that active carbon adsorbed no alkali urate from a solution whose pH value varied from 6 to 10, and concluded that the oxidation is purely a contact process. The increase in rate of reaction observed here can be attributed to the larger surface area per unit weight of carbon that is produced as the temperature of activation is increased. This is quite different from the decomposition of hydrogen peroxide, in which reaction the type of oxide present on the surface is the factor which decides whether or not a particular sample of carbon is to be an active catalyst. SUMMARY

The activity of activated carbon was measured by its effect on the rate of decomposition of hydrogen peroxide, and on the rate of autoxidation


ACTIVATED CARBON A 8 A CATALYST

85

of stannous chloride and potassium urate. The effect of temperature of activation on the amount of iodine adsorbed from aqueous solutions was determined. The effect on the rate of decomposition of hydrogen peroxide when catalyzed by activated carbon was determined for the temperature of activation, source of carbon, initial peroxide concentration, variation of both quantity of catalyst and of peroxide, temperature of reaction, pH, metallic salts, and repeated use of the same catalyst. The decay in activity of the carbon during peroxide decomposition was discussed from the point of view of surface oxides of carbon and the chain mechanism of peroxide decomposition. REFERENCES STYLE,D. W. G . : J. Chem. SOC.1030,598608. V. L.:J. Phys. Chem. 26, 19 (1921). (2) BOHNSON, W.,AND FIRTH,J. B . : J. Phys. Chem. 28, 1136 (1924). (3) FARMER, (4) FILSON, G.W., AND WALTON,J. H . : J . Phys. Chem. 36,740 (1932). (5) FIRTH,J. B., AND WATSON,F. S.: J. Chem. SOC.123, 1750 (1923). (6) FIRTH,J. B., AND WATSON,F. S.: J. Phys. Chem. 20, 987 (1925). M.:Compt. rend. 101, 949 (1930). (7) FR~REJACQUE, (8) HABER,F . , AND WEISS,J.: Naturwissenschaften 20,948 (1932);Proc. Roy. SOC. (London) A147, 332 (1934). (9) HABER,F., AND WILLSTHTTER, R.: Ber. 64, 2844 (1931). (10) HARING, R. C., AND WALTON, J. H . : J . Phys. Chem. 37,375 (1933). (11) KING,A.: J. Chem. SOC.1036, 1688-92. (12) KING,A.: J. Chem. SOC.1038, 991-7. C. G . : Kolloid-Z. 60, 21 (1934). (13) KING,A,, AND LAWSON, E.: Rec. trav. chim. 46, 453 (1927). (14) KISS,A., AND LEDERER, I. M.:J. Am. Chem. SOC.64,4473 (1932). (15) KOLTHOFF, (16) KRUYT,H . R., AND DE KADT,G. S.: Kolloid-Beihefte 32, 249 (1931). G.:Compt. rend. 144, 357 (1907). (17) LEMOINE, D.A.: J. Am. Chem. SOC.36, 878 (1914). (18) MACINNES, E.J.: J. Phys. Chem. 36, 2967 (1932). (19) MILLER, N.,SCHATUNOWSKAJA, H., AND TSCHMUTOW, K . : Z. physik. Chem. (20) SCHILOW, A140, 211 (1930). (21) TRTJSZKOWSKI, R . : Biochem. J. 24, 1349 (1930). (22)WEISS,J.: J . Phys. Chem. 41, 1107 (1937). (23) WEISS,J.: Trans. Faraday SOC.31,1547 (1935). (1) ALLMAND, A. J.,

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