ACTIVATED NITROGENOUS CARBONS
133
they have been interpreted qualitatively. The polarimetric data appear to be inconsistent with the equilibrium constants derived by Levy. REFERENCES (1) FRIEDEN, E. H., DUNN,M. S., (2) FRIEDEN, E. H., DUNN,M, S.,
(3) (4)
AND CORYELL, C. D . : J. Phys. Chem. 48,216 (1942). AND CORYELL, C. I),:J. Phys. Chem. 47.10 (1943). FRIEDEN, E. H., DUNN,M. S., AND CORYELL, C. b.:J. Phys. Chem. 41, 20 (1943). FRIEDEN, E. H., DUNN,M. S., AND CORYELL, C. D . : J. Phys. Chem. 4 7 , s (1943).
(5) LEVY,M.: J. Biol. Chem. 109, 365 (1935).
THE CATALYTIC ACTIVITY OF ACTIVATED NITROGENOUS CARBONS' PAUL F . BENTE
AND
JAMES H. WALTON
Uepartment of Chemistry, University of Wisconsin, iMadison, Wisconsin
Received M a y 80, I$&
It has been shown by Larsen and Walton (10) that activated carbon prepared from ash-free gelatin is much more active in catalytically decomposing hydrogen peroxide than carbon derived from sugars. It was the object of the present investigation to prepare activated nitrogenous carbons from different sources and to study the catalytic activities of these carbons upon (1) the decomposition of hydrogen peroxide, (8) the oxidation of hydroquinone, and (3)the oxidation of potassium urate. PREPARATION OF CARBONS
The method of preparing the carbon was essentially that used by Larsen and Walton and consisted in charring the material at about 450°C., grinding to 100 mesh or more, heating in a vacuum a t 1OOO"C. for 45 min., and then activating the residue by heating in moist oxygen a t 875°C. for 12 hr. Table 1 lists the sources of carbon and gives the ash and nitrogen content. The yield of activated c.arbon was 20 to 50 per cent of the char. The ash content was not considered sufficiently high to warrant any purification. Spectroscopic examinations of the ash indicated that traces of many metals were present. Iron was pkesent to an insignificant extent; this is interesting in view of the fact that the claim has been made that for certain reactions the catalytic activity of carbons is due to the presence of iron-carbon and iron-carbon-nitrogen complexes (14, 19). The per cent nitrogen present in the carbons as nitrogen of constitution was estimated by Kjeldahl analysis. Determinations of nitrogen by the Dumas 1 This investigation WIW financed by a grant from the Research Committee of the University of Wisconsin, Dean E. B. Fred, Chairman.
134
PAGL F. BEXTE AND JAMES H. WALTON
method were not as satisfactory, because it \vas difficult to remove all the adsorbed nitrogen from the carbons. It is not likely that the nitrogen of constitution of the carbons is adsorbed ammonia, since the carbons were heated in a vacuum a t 1000°C. before activating them at 875°C. Unless otherxise indicated, the above carbons \yere used in all the expeiiments to be described. Other carbon samples are marked B and C. TABLE 1 Yilrogenous carbons used SOURCE OF CAPBON ~
I
Hexamethylenetetramine. . . . . . . . . . . . . . . . . . . . . . . . Gelatin . . . . . . . . .. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Glucosazone, . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ’ Lactose (90%) and hexamethylenetetramine (1O%)i Lactose (900Jo) and urea (IO’%). . . . . . . . . . . . . . . . . . Lactose., . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ,
per ceni
0.5-1.0 19Ccn‘
20.5 26.5 28 11
0.25 0.18 0.11* 0.02 0.06 0.22
I ~
per cent
3.0 ,
i
~
3.6 3.9 2.4
2.9 None
* Calculated from per cent ash in activated carbon.
FIG.1 . The effect of activation temperature on the time of half-life of decomposition of hydrogen peroxide by carbon. Curve 1,O.lOO g. of gelatin carbon; curve 2,0.050 g. of gelatin carbon; curve 3, 0.300 g. of 90 per cent lactose and 10 per cent hexamethylenetetramine carbon. T H E CdTALYTIC DECOMPOSITION O F HYDROGEN PEROXIDE
Temperature of activation of carbons To determine the optimum conditions for preparing highly active catalytic carbons, two samples, one prepared from gelatin, the other from a mixture of 90 per cent lactose and 10 per cent hexamethylenetetramine, were activated at various temperatures. The catalytic activity of each carbon was measured by
135
ACTIVATED NITROGENOUS CARBONS
determining the rate-at which it decomposed hydrogen peroxide (5, 10). ,411 experiments were carried out a t 25"C., using 15-ml. portions of solutions containing 50 ml. of available oxygen. All gas volumes have been corrected to standard conditions. The conditions for optimum activation are best shown graphically (figure 1). Since the carbons differ greatly in activity, two scales have been used. Activation a t temperatures below 700°C. produced very inactive carbons, not shown in the figure, When a smaller amount of a particular carbon catalyst was used, the optimum temperature of activation could be more accurately observed. Consequently 875°C. was the activation temperature used for the carbons prepared in the folloiving experiments. This value compares favorably with that of 825475°C. noted by Larsen and Walton (10) for sugar carbons.
Catalytic activities of carbons The catalytic activities of the various carbons were compared as in the study of the temperature of activation. Duplicate experiments for a given sample of TABLE 2 Decomposition of hydrogen perozide by uarioua carbons TIME OF EAIF-LIFE
SOURCE OF CARBON
WEIGET USED
March, 1940 grams
Hexamethylenetetramine ......................... Gelatin. ......................................... Glucosazone ..................................... Lactose (90%) and urea (10%). . . . . . . . . . . . . . . . . .J Lactose (90%) and hexamethylenetetramine (10%) Lactose. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
0.025 0.050 0.050 0.050
July, 1941
secadr
20 256
32 266 759 2377 690 781 3160
540
0.100 0.100 0.300
1
sccadr
2.945
carbon agreed within 2 per cent. The results are given in table 2. From this table it is evident that the nitrogenous carbons are much more active than the lactose carbon. Even within the series of nitrogenous carbons, the range of activities is surprisingly large. The great activity of the hexamethylenetetramine carbon is especially noteworthy.
The decay factor These experiments also show that the carbons have decreased in catalytic activity after standing for 15 months. Though the samples were kept in evacuated desiccators, they came into contact with air when the desiccators were opened. Since Larsen and Walton showed that moist oxygen readily deactivates carbon, this decay is not surprising. It has been assumed that the moist oxygen is also responsible for the deactivation of the carbon during the decomposition of a solution of hydrogen peroxide. The very active nitrogenous carbons make it easily possible to follow this reaction to completion, thereby allowing any variation in the volume of oxygen
136
PAUL F. BENTE AND JAMES
E. WALTON
evolved to be measured accurately. This has not been feasible with the less active carbons prepared by previous investigators. The amount of oxygen obtained must be corrected for gases released on wetting the carbon (said by McBain to be nitrogen (11)). Using 0.100-g. samples of carbon C, prepared from a mixture of 90 per cent lactose and 10 per cent hexamethylenetetramine, it was found that when the reaction was complete 0.40ml. of oxygen was adsorbed by the carbon, since only 49.60m1. of the 50.00 ml. available wasevolved. Corrections meremade for blanks, using water instead of hydrogen peroxide. The time of half-life was 331 sec. at 25OC. When the reaction was carried out a t 15"C., the carbon adsorbed 1.35 ml. of oxygen, and the time of half-life increased to 481 sec. When the carbon was merely suspended in water and agitated in the presence of oxygen, it also adsorbed this gas. Furthermore, the carbon so treated showed TABLE 3 Temperature eoefieients and inhibitory effect of potassium cyanide for the decomposition of hvdroaen peroxide bv carbons TlhZ OF EA=-LIBE
SOURCE OF CARBON
WEIGET USED
25'C.
Lactose. ......................... Hexamethylenetetramine (B). , , . . Gelatin (B). . . . . . . . . . . . . . . . . . . . . . Lactose (90%) and hexamethylenetetramine (10%). . . . . . . . . . . . Lactose (goo/,) and urea (10%). . . . Lactose (90%) and hexamethylenetetramine (10%) (B). . . . . . . .
___
PEWERATUBE
15'C.
0.001 N X N , 2S"C
__ -
COEFFICIENT
pams
7?liilUtCS
minutes
minufes
0.300
97.5 0.63 10.2
58.0 1.60 14.0
1.44
0.025 0.050
67.5 0.45 7.4
0.100 0.100
12.6 8.1
16.6 9.8
27.3 16.8
1.32 1.21
0.100
4.7
5.6
11.3
1.19
1.40 1.38
that its activity had decayed. Thus a sample of carbon which had adsorbed 1.25 ml. of oxygen by agitation for 19 hr. showed a half-life increase of about 100 per cent. However, a carbon which had decomposed hydrogen peroxide and adsorbed about 0.40 ml. of oxygen, when used a second time showed greater decay, the period of half-life having been increased by 400 per cent. The decay in activity may be attributed to the chemisorption of oxygen by the carbon to form less active surface oxides. That oxygen derived from hydrogen peroxide causes greater decay than molecular oxygen may be explained by assuming that the hydrogen peroxide oxygen attacks the active centers where the decomposition occurs, while the molecular oxygen does not preferentially poison these centers. Since chain mechanisms for the decomposition are involved, this effect is very noticeable. From the reciprocal of the half-lives of the reaction the temperature coefficient is found to be 1.45. However, it should be noted that the temperature coefficient is very dependent on the decay factor, which may vary for different carbons.
ACTIVATED NITROQENOUS CARBONS
137
If the carbon picks up more oxygen a t the lower temperatures and this oxygen causes greater decay, the reaction will naturally be slower, the time of half-life greater, and the temperature coefficient larger. This in part may account for the fact that Fowler and Walton ( 6 ) observed the temperature coefficient to be only slightly more than 1, while Larsen and Walton found it to be about 1.90. Even for the uniformly prepared carbons studied here, the temperature coefficient was found to vary for different preparations (see table 3).
Inhibition by potassium cyanide Warburg (19) found that low concentrations of cyanide inhibited certain oxidation reactions in which a nitrogenous carbon having metallic impurities was employed as catalyst. Skumburdis (16), using potassium cyanide of 1 and 10 per cent concentrations, found that the decomposition of hydrogen peroxide by carbons was very greatly promoted. Since the effect of equivalent concentrations of hydrogen cyanide was definitely inhibitory, the promotion was explained as due to the free alkali formed by hydrolysis. It was of interest therefore to determine the effect of low concentrations of potassium cyanide on this reaction, using the types of carbons described above. The results of experiments with freshly prepared hydrogen peroxide solutions containing 0.001 N potassium cyanide are included in table 3. For the different nitrogenous carbons the presence of the potassium cyanide approximately doubled the time of half-life for the decomposition. However, for the lactose carbon little effect was noticed. The nitrogenous carbons apparently contain certain active centers which are responsible for their enormous catalytic activity. In addition, other parts of the surface contribute to the activity of the carbon. This is evidenced by the effect of the presence of potassium cyanide, which inhibits the action of the nitrogenous carbons but has little effect on the lactose carbon. The experimental evidence presented supports the oxide theory often used for active carbons. According to this theory the properties of active carbons depend largely upon the existence of two types of surface oxides, one formed a t high, the other a t low temperatures. In these carbons the high-temperature oxide, best formed a t 875"C., catalyzes the decomposition of hydrogen peroxide. The decay effect observed when the carbon picks up oxygen shows that such oxide centers are capable of being partly destroyed. That such a small amount of oxygen can have such a large effect indicates that a great deal of the surface may be relatively inactive to begin with. The poisoning experiments cited above further indicate that these highly active centers probably involve the presence of metallic impurities and nitrogen, sinw the cyanide has been shown by Warburg to be selectively adsorbed on such surface complexes of carbon (19). In this connection it is interesting to note that for the nitrogenous carbons the order of decreasing catalytic activity approximates the order of dxreasing ash content (cf. table 1). Since the ash analyzes to show that many metallic constituents are present, it is difficult to enlarge upon this.
138
PAUL F. BENTE AND JBhlES K. WALTON THE OXIDATION O F HYDROQUINOXE
The catalytic effect of carbon in the autoxidation of solutions of hydroquinone is recorded by Matsui (12), Gandini @), and others. Gandini states that in ether solution Norite rapidly catalyzes the reaction to form quinhydrone, which in turn is slowly oxidized to quinone. This reaction was used by the authors to compare the various catalytic activities of the nitrogenous carbons and not to determine its exact nature. Most of the carbons used were too inactive to carry the oxidation as far as the quinhydrone point. Hovever, in the case of the highly active hexamethylenetetramine carbon the reaction proceeded as observed by Gandini.
ACTIVATIONTEMPERATUeE
‘c
FIG.2. The effect of activation temperature on the carbon-catalyzed oxidation of hydroquinone i n alcohol. 15-ml. samples of 0.30 molar hydroquinone. 0.100-g.samples of gelatin carbon.
Apparatus and technique The same apparatus and technique used for the decomposition of hydrogen peroxide were employed, except that absorption instead of evolution of oxygen was now measured.
Effect of temperature of activation To determine the optimum temperature for producing an active carbon catalyst for this reaction, 0.100-g. samples of gelatin carbon activated a t six different temperatures were treated with 15 ml. of 0.300 molar hydroquinonc in oxygen-free 95 per cent ethyl alcohol. The results of the experiments (figure 2) show that the gelatin carbon is most active when activated at 875°C. This differs from the data of King (9), who for sucrose carbon obtained q o s t activity a t 420°C. activation and least at 800-900°C. The activity of this 875°C.gelatin carbon for this oxidation reaction is much greater than that of King’s sucrose carbon.
139
ACTIVATED NITROGENOUS CARBONS
TABLE 4 Effect of solvent on catalytic oxidation of hydroquinone
I SOLVENT
I
OXYOEN ADSORBED IN 30 MIN.
0.2 K. of 840°C.
lactose cnrbon
1
O . Z g . , o f 840°C. gelatin carbon
0.1 E. of 875°C.
gelatin ' carbon (C) ml.
m!.
m!.
Water. . . . . . . . . . . . . . . . . . . . . . . . . Ethyl alcohol (95%) . . . . . . . . . . . . Acetone. . . . . . . . . . . . . . . . . . . . . . . Ethyl ether., . . . . . . . . . . . . . . . . . . Dioxane. ......................
1
0.55
16.65 9.75
2.80 2.10 1.40 0.15 0.13
0.10 0.05 0.05
1.65 0.60*
* Amount of oxygen absorbed in 70 min., using 0.2 g. of carbon.
5
IO
IS
TIME IN MINUTES FIG.3. The oxidation of 15-ml. samples of 0.50 molar hydroquinone (aqueous), using hexamethylenetetramine carbon as catalyst,
Solvent effect To determine the most favorable solvent for the reaction, fixed amounts of carbon were added to 15-ml. portions of 0.3 molar oxygen-free solutions. The greateRt activity was noted in water solutions, followed in order by 95 per cent
140
PAUL F. BENTE AND JAMES R . WALTON
ethyl alcohol, acetone, ethyl ether, and dioxane. Dioxane almost completely suppressed oxidation. The data are given in table 4. The decrease in activity with change of solvent does not correspond to the changes in oxygen solubility, though it does agree roughly with changes i n the dielectric constants. Other factors which are involved but which were not determined are the relative adsorptive capacities of the carbons towards hydroquinone, quinhydrone, and quinone in the various solvents. Also, in the experiments with water the oxidation products, being less soluble, may precipitate at much lower concentrations than in the organic solvents. From figure 3 it is evident that the products, though insoluble, must be desorbed for the catalyst does not lose its activity, the reaction rate remaining constant as the oxidation proceeds. In the aqueous medium the oxidation actually proceeds somewhat faster in the more dilute solutions containing hexamethylenetetramine carbon. Since the solubility of oxygen in 0.1 and 0.5 molar solutions of hydroquinone was found to TABLE 5 Catalytic activities of the various carbons -
SOURCE
or
CAXBON
i
WEIOHT USED
1
___ ~
15 14
Gelatin, . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Glucosaaone . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Lactose (90%) and hexamethylenetetramine (10%)
0.103 0.1O0 0.100
OXYOEN
TIP€
60 60 60 60 60
ABSOQRED
34.00 41.60
9.60 1.05
0.70 0.70 0.40
be the same, the change in rate of reaction with change in concentration cannot be explained as due to different oxygen solubilities.
Catalytic activities of the carbons Using 15-mi. samples of aqueous 0.50 molar hydroquinone, the catalytic activities of the various carbons activated at 875OC. were compared. The results are shown in table 5. When these catalytic activities are compared with those found for the decomposition of hydrogen peroxide, it is evident that the order of the decreasing activities of the carbons is the same.
Characteristics and inhibition of the reaction The oxidation of 15-ml. samples of 0.50 molar hydroquinone catalyzed by hexamethylenetetramine carbon activated at 87OoC. has zero-order characteristics, as shown by figure 3. This fact is useful in making extrapolations to zero time to determine the necessary corrections for gases desorbed on wetting the carbon. Figure 3 also shows that the smaller the ambunt of carbon used the more the reaction departs from zero order. If less active carbons are used the
ACTIVATED NITROGENOUS CARBONS
141
reaction has no definite order, but is complex. The same is true when the hydroquinone concentration is reduced. Thus the zero-order characteristics indicate the extent to which this reaction is dependent on factors such as the rate of diffusion of the reactants to the catalyst. If the hexamethylenetetramine carbon is previously suspended in water and exposed to oxygen for some 20 hr., it shows no catalytic decay as in the hydrogen peroxide experiments.
TIME in MINUTEJ FIG.4. The effect of 0.001 iV potassium cyanide on the catalytic oxidation of 0.30 molar hydroquinone (aqueous). Curves I and 11, 0.100 g. of hexamethylenetetramine carbon B; curves I11 and IV, 0.300 g. of 90 per oent lactose and 10 per cent hexamethylenetetramine carbon B; curves V and VI, 0.500 g. of lactose carbon.
The inhibitory effect of 0.001 N potassium cyanide on 0.5 molar hydroquinone was also investigated. Some of the data are given in figure 4. For the highly active carbon the reaction was very markedly inhibited (cf. curves I and 11, figure 4), while for the relatively inactive carbons including the lactose carbon the reaction was actually promoted (cf. curves V and VI). A carbon selected for its intermediate activity showed promotion for the early part of the reaction and inhibition thereafter (cf. curves I11 and IV).
142
PAUL F. UENTE A S D JAMES H. WALTOS
Since the potassium cyanide hydrolyzes to give free alkali and since the oxidation of hydroquinone in alkaline medium is very rapid, the promoting effect of the potassium cyanide in the presence of relatively inactive carbons is probably due to the presence of free alkali. When a more active catalyst is used this effect may still be observed, but it is more likely to be masked completely by the inhibition caused by selective adsorption of the cyanide on the catalytically active centers. The effect of activation temperature in preparing the carbons indicates that for nitrogenous carbons a temperature of 875"C., where the high-temperature oxides are supposedly best formed, is the optimum for forming active centers with proper spacing. Thus the promotion effect due to the presence of both ash and nitrogen is again evident from the potassium cyanide inhibition. However, exposure of the carbon to oxygen does not modify the activity as for the decomposition of hydrogen peroxide, indicating that for the hydroquinone oxidation it is not the type of surface oxide that matters but some more fundamental factor, such as proper spacing of active points, caused by the presence of tjhenitrogen and ash during formation of the oxide surface a t 875°C. THE OXIDdTION OF POTASSIUM URATE
The oxidation of an alkaline solution of potassium urate in the presence of activated carbon has been studied by Frerejacque ( 7 ) ,Truszkowski (18), Zylbertal (20),and Larsen and Walt,on (10). The reaction, which requires 2 moles of oxygen to react with 3 moles of potassium urate, is zero order. The method of following this reaction was the same as that used for the oxidation of hydroquinone. In all experiments 15-ml. samples of an oxygen-free solution containing 12.5 g. of uric acid and 15.6 g. of potassium hydroxide per liter were treated ivith the various carbons and the rate of oxygen absorption measured.
Catalytic activities The zero-order characteristics of this reaction are very pronounced for the highly active hexamethylenetetramine and gelatin carbons. However, for less active carbons or for smaller amounts of more active carbons the reaction departs from zero order and becomes complex, as in t'he case of the hydroquinone reaction. Since the senior author (10) noted that the activity of sugar carbons increased with activation temperature, experiments using the highly active carbons activated at 875°C. were carried out. The data are given in table G . Table 6 shows that the carbons arrange themselves in the same order of decreasing catalytic activities as observed for the other reactions studied. Inhibition Representative results of experiments carried out to determine the inhibition due t o previous exposure of the carbon to oxygen and the presence of potassium cyanide are given in figure 5 . Curves I and I1 show that a hexamethylenetetramine carbon when exposed to water and oxygen, as described previously, showed a definite catalytic decay. Furthermore, the inhibitory effect of 0.001
I43
BCTIV.ITED NITROGESOUS CARBOSS
N potas-sium cyanide was evident for all the nitrogenous carbons, being most pronounced for the highly active hexamethylenetetramine carbon (cj. curves I and 111). For the lactose carbon, however, the potassium cyanide had no TABLE 6
The oxidation of potassium urate SOURCE OF CARBON
~
Hexamethylenetetramine. ........................ Gelatin. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Glucosazone . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Lactose (90%) and hexamethylenetetramine (10%) Lactose (90%) and urea (10%). . . . . . . . . . . . . . . . . . . Lactose.. ........................................ None. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
WEIGHTUSED
1
z:iE&1I,
grams
ml.
0.100 0.200
10.50 10.60 10.20 10.10 10.20 4.35 0.85
0.200 0.200 0.200
0.200
TlYE
miniiles
1.25 6 8
30 30 30
30
TIME IN MINUTES FIG.5 . The effect of 0.001 .V potassium cyanide on the catalytic ohidation of potassium urate. Curves I , 11, and 111, 0.100 g. of hexamethylenetetramine carbon B; curves IV and V, 0.100 g. of glucosazone carbon; curves VI and VII, 0.300 g. of lactose carbon.
effect (curves VI and VII), showing that the reaction on the normal carbon surface mas unaffected by the presence of the poison. The experiments for this reaction again indicate that the more active nitrogenous carbons owe their catalytic activities t o active centers which have both nitrogen and ash present as promoting agents and which may be selectively
144
PAUL F. BENTE AND JAMES H. WALTOS
poisoned by adsorbing the cyanide. Here the type of oxide may also be of significance, since decay occurs on exposure to oxygen. Thus the mechanism of the oxidation of alkaline potassium urate probably differs from that for the oxidation of hydroquinone, since for the latter reaction the carbon shows no decay on exposure to oxygen. DETERMINATIONS OF SGRFACE ARE4
Effect of nitrogen on subdivision of carbons All the nitrogenous carbons studied in this laboratory are considerably more active as catalysts than the carbons prepared from non-nitrogenous materials. It is generally agreed (13, 14) that the presence of nitrogen in organic compounds produces carbons with relatively larger surfaces. From photomicrographs (110 magnification) of the six carbons studied it is evident that the nitrogenous carbons had undergone considerably more rupturing than the lactose carbon. Figure 6 gives an illustration of this. When the photomicrographs of the carbons are arranged in the order of their decreasing particle sizes, one obtains approximately the same order as that observed for the decreasing catalytic activities. The catalytic properties studied often differed so greatly that moderate differences in surface area can hardly account for these variations. Further proof that nitrogenous carbons in general have more surface than sugar carbons is shown by the measurement of their bulk densities, as described by McBain (11). Table 7 shows that nitrogenous carbons are more dispersed, but there is no correlation bekeen bulk densities and catalytic activities. Absolute surface areas, ete. Since more accuratr observations were desirable, the absolute surface areas weie determined by the gas-adsorption method of Emmett and Brunauer (4). Surface areas mere calculated from the amount of nitrogen necessary to cover the carbon surface with a monolayer of gaq. An adsorption apparatus of Pyres resembling that of Benton and White (2) vas used. Before measuring the admm. pressure for 16 hr. a t sorption, each sample of carbon n’as evacuated at a temperature of 360-400°C. The surface areas are given above in table 7 , one column being calculated from the slope of a straight-line function for the isotherm described by Emmett, the other from the observed saturation point found on the normal isotherm. The results show that there is no direct correlation betmen these absolute surface areas and the catalytic activities observed. Experiments on adsorption of iodine were also carried out on the various carbons. The milligrams of iodine adsorbed per gram of carbon are given in table 7. It is significant that the ratio of absolute surface area t o the amount of iodine adsorbed is fairly constant, best agreement being obtained nith surface area values calculated from the observed saturation value of nitrogen on the carbon. Thus for these carbons each milligram of iodine absorbed is equivalent to about 10 square meters of absolute surface area, as found by the lovv-tempera-
ACTIVATED NITROGENOUS CARBONS
145
146
PAUL F. BEXTE A K D JhhlES H. WALTOK
ture nitrogen-adsorption method. After this work was begun, Smith, Thornhill, and Bray (17) showed that for carbon blacks absolute surface areas are proportional to iodine adsorption. In their work they found that 1 mg. of iodine is equivalent to 0.77 square meter of absolute surface area, a value which is in keeping with the possibility of the formation of a monolayer of adsorbed iodine. If the adsorbed iodine does form a complete monolayer on the carbon, then the active carbons used here must have nearly 90 per cent of their absolute surface area in micropores which are inaccessible to the iodine molecule. Such a value is probably too high and indicates rather that the type of surface as well as the available surface determines the amount of adsorption. In support of this mag be cited the evidence of Larsen and Walton (lo), shobving that adsorptive capacity towards iodine decreases materially if the carbon is allowed to stand overnight in contact with hydrogen peroxide, a condition which brings about a change in the surface oxide but which cannot change the surface area materially. TABLE 7 Surface properties of the carbons BSOLUlX SURFACE AREA SOURCE OF CARBON
1
Hexamethylenetetramine. . . . . . Gelatin . . . . . . . . . . . . . . . . . . . . Glucosssone . . . . . . . . . . . . . . . . Lactose (90%) and hexamethyl-1 enctetrsmine (loo/,),. . . . . . ' Lactose (90%) and urea (10%). Lactose . . . . . . . . . . . . . . . . . . . .
I
~
RATIO OF COLUMN 3
BULK DENSITY
IO COLUVS 4
5.7.
m.
gram.?per cc.
per gram
0.103 0,590 0.442
2080 575 060
10.9 10.5
0.596 0.562 0.703
802 907 896
9.2 0.2 9.5
9.9
896
i
94
Thus one may conclude that, though iodine-adsorption values are directly proportional to absolute surface areas, the proportion holds only when a comparison of similarly prepared carbons is made. Once the proportion for a given type of carbon has been determined, the absolute areas may be calculated approximately from the iodine-adsorption values, but only for carbons of that type. The gelatin carbon is of special interest, since it has very high catalytic activity for the three reactions studied but by far the lowest adsorptive capacity towards iodine in solution and nitrogen gas a t loiv temperatures. This same carbon, when tested for phenol adsorption (l),was found to be very inferior t o a commercial carbon used in water purification. However, the commercial carbon was much less effective than the gelatin carbon for decomposing hydrogen peroxide solutions. If the various carbons had the same number of catalytically active points of the same kind in a given area, one might logically expect the catalytic activities to be directly proportional to the surface areas. That no such correlation exists is not
ACTIVATED NITROGEKOC'S CARBOKS
147
very surprising, since the carbons vary considerably in both ash and nitrogen content, both of which are believed to have pronounced promotion effects. It is more surprising that the carbons show the same order of catalytic activities for all three reactions. SUMMARY
The catalytic activities of several nitrogenous carbons and one sugar carbon mere compared by measuring their effects on the rate of decomposition of hydrogen peroxide, the rate of oxidation of hydroquinone in various solvents, and the rate of oxidation of alkaline potassium urate. The optimum temperature of activation was found to be 875°C. The carbons can be arranged in the same order of decreasing activity for all three reactions. The carbon prepared from hexamethylenetetramine showed exceptional catalytic activity. The decay factor in the decomposition of hydrogen peroxide is shown to b'e due to adsorption of oxygen on the active centers and is explained on the basis of the oxide theory. The possible effect of such decay on the temperature coefficient is discussed. Decay on exposure to oxygen was also noted for the oxidation of potassium urate but not for the oxidation of hydroquinone. Potassium cyanide (0.001 N ) inhibits the catalytic action of the nitrogenous carbons in all three reactions. It does not affect the catalytic properties of lactose carbon. The inhibitory effect indicates that a small per cent of the total surface of the nitrogenous carbon is responsible for most of the catalytic activity. Photomicrographs of the carbons were prepared and the bulk densities determined. The absolute surface areas measured by the low-temperature gasadsorption method were found to be proportional to the iodine adsorptive capacities of the carbons. No direct correlation was found to exist between these sorptive properties and the catalytic properties of the hrbons. REFERENCES (1) AMERICAN PUBLICHEALTHASSOCIATION: Standard Methods of Water Analysis, 8th edition. New York (1936). (2) BENTON, A. F., AND WHITE,T . A , : J. Am. Chem. Soo. 62, 2325 (1930). (3) BERKMAN, S., MORRELL, J. C., AND EGLOFF,G.: Catalysis, Inorganic and Organic, p. 41. Reinhold Publishing Corporation, New York (1940). (4) BRUNAUER, S.,EMMETT, P. H., AND TELLER, E . : J. Am. Chem. Soc. 80,309 (1938). (5) FILSON, G. W., AND WALTON, J. H . : J. Phys. Chem. 36, 740 (1932). (6) FOWLER, D . , AND WALTON, J. H . : Rec. trav. chim. 64,476 (1935). (7) FREREJACQUE, M . : Compt. rend. 191, 949 (1930). (8) GANDINI, A.: Gam. chim. ital. 63, 9 (1933). (9) KING,A . : J. Chem. Soc. 1936, 1688. (10) LARSEN, E. C., AND WALTON, J. H . : J. Phys. Chem. 44,70 (1940). (11) MCBAIN,J. W . : The Sorption of Gases and V a p o u w b y Solids, pp. 79, 93. George Routledge and Sons, Ltd., London (1932). (12) MATSUI,M . : Mem. Coll. Sei. Eng., Kyoto Imp. Univ. 1,386 (1909). F., AND RADU,A , : Ber. 67, 1221 (1924). (13) PANETH, (14) RIDEAL, E. K., AND WRIGHT, W. hf.: J. Chem. SOC.1926, 1813. (15) SCHWAB, G.-M., TAYLOR, H. S., AND SPENCE, R . : Catal@is, p. 249. D. Van Nostrand Co., Inc., New York (1937).
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(16) SKUMBURDIS, K.:Kolloid-2. 66, 156 (1931). (17) S w n i , W. R., THORNHILL, F. S., AXD BRAY,R. I.: Ind. Eng. Chem. 83, 1313 (1941). (18) TRCSZKOWSKI, R.:Biochem. J. 24, 1349 (1930). , Biochem. Z.ll9.134 (1921);138,266 (1923);146,461 (1924). (19) ~ ~ A R B U R G0.: (20) ZYLBERTAL, S.: Biochem. 2. 236, 131 (1931).
STUDIES O S AGISG AND COPRECIPITATION. XXXVII
THE DISTRIBUTIOX COEFFICIENT OF ARSENATE BETWEEN xfZIAGNESIUM AMMONIUM PHOSPH.4TE AND SOLUTION’ I. M. KOLTHOFF
AND
C.W. CARR
School of Chemistry, Institute of Technology, Universilyof Minnesota, Minneapolis, Minnesota Received November 20, 1942
The orthorhombic salts MgNH4P04 6H20 and Mgirr”dsO4.6H20 are isomorphous. Their axial ratios are almost identical, being as follows (2): LlgNHrPOa. 6H20 MgNH&O4*6HpO
a : b : c = 0.5667:1.0:0.9122 u : ~ : =c 0.5675: 1.0:0.9122
Analytical use of the mixed-crystal formation between the phosphate and the arsenate is made in the quantitative coprecipitation of traces of dissolved arsenate with magnesium ammonium phosphate (1). The arsenate in the precipitate is determined by classical procedures. In this way as little as 0.075 mg. of arsenic dissolved in 500 ml. of solution could be determined with an accuracy of 2 per cent. The quantitative coprecipitation is accomplished by adding 500 mg. of phosphorus pentoxide in the form of monopotassium phosphate, 1 ml. of hydrochloric acid, and 10 ml. of magnesia mixture2 to 500 ml. of solution. The solution is neutralized with ammonia and after most of the precipitate has been formed, 5 ml. of concentrated ammonia is added. The precipitate is filtered after 2 to 4 hr. of standing and washed with dilute ammonia. MECHAXISM OF THE COPRECIPITATIOh’
The quantitative precipitation of small amounts of arsenate in the absence of phosphate requires a long time. This precipitation from supersaturated solutions is promoted by the addition of magnesium ammonium phosphate. Even fairly perfect crystals of magnesium ammonium phosphate which are not subject 1 This paper is based upon a thesis submitted by C. W. Carr t o the Graduate School of the University of Minnesota in partial fulfillment of the requirements for the degree of Master of Science, June, 1939. 2 Fifty grams of MgClz 6H10 and 100 ml. of water. A slight excess of ammonia was added and the solution allowed t o stand overnight. If a precipitate waa formed, the solution was filtered from it. It mas then slightly acidified with hydrochloric wid and diluted t o 1 liter.
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