A Comparison of the Catalytic and Chemical Characteristics of Cubic

A COMPARISON O F T H E CATALYTIC AND CHEMICAL CHARACTERISTICS OF CUBIC AND RHOMBOHEDRAL Fe203* B Y P . H. EMMETT AND KATHARINE S. L O V E

Introduction For many years it has been known that a magnetic form of ferric oxide exists. As pointed out by Herroun and Wilson’ such an oxide was mentioned ~ later by both Malaguti3 and Livemidge.‘ A review by Robbins in 1 8 5 9 and of the literature5 in more recent years reveals many additional references to a magnetic Fe203. Welo and Baudisch6 have shown that “magnetic Fe2O3’’ is not merely a thin layer of non-magnetic FezO3 over magnetic Fe304. They have pointed out the additional interesting fact that the x-ray powder photographs of this magnetic Fez03are indistinguishable from those of Fe304. Hendricks and Albrecht’ working with CozO3zFe2O3and CoOFez03,compounds that are analogous to the above-mentioned cubic Fe203and Fe304, concluded that in this case also, within the error of experimental observation, no difference exists between the powder photographs obtained from the two forms. The following conclusions can be drawn from the previous work: ( I ) within the limits of error of present observations the x-ray powder photographs obtained from cubic and from rhombohedral Fez03are identical; ( 2 ) the magnetic permeability of the cubic Fe203is similar to that of Fe304;( 3 ) the cubic modification of the oxide can be changed into the more stable rhombohedral modification by heating to various temperatures between j 50’ and IOOO’C.The transition temperature depends upon the method of preparation of and the impurities contained in the sample. Welo and Baudisch* studied the comparative chemical and catalytic characteristics of the two oxides. They prepared samples of cubic Fe203 by heating hydratedg Fe304for several hours in a stream of air at a tempera*Fertilizer and Fixed Xtrogen Investigations, Bureau of Chemistry and Soils, Washington, D. C. l Herroun and Wilson: Proc. Phys. SOC.London, 41, (pt. 11) IOO ( 1 9 2 8 ) . * Robbins: Chem. Kews, 1, 1 1 (1859). Malaguti: Ann. Chim. Phys. (3), 69, 2 1 4 . Liversidge: Trans. Aust. Assoc. Hobart Meeting ( 1 8 9 2 ) . Hauser: Ber., 4 0 , 1958-60 (1907); Sosman and Hostetter: Bull. Am. Inst. Mining Eng. 1917, 907-31; Abraham and Planiol: Compt. rend., 180, 1 3 2 8 9 ( 1 9 2 5 ) ; Chevallier: 180, 1473-5 (1925); 184, 674-6 ( 1 9 2 7 ) ; Sosman and Posnjak: J. Wash. .kcad. Sci., 15, 329-42 (1925); Wedekind and Albrecht: Ber., 59, 1726-30 (1926); 60, 2239-43 ( 1 9 2 7 ) ; Huggett and Chaudron: Compt. rend., 186, 1 6 1 7 9 ( 1 9 2 8 ) . Welo and Baudisch: Phil. Mag., 50, 399-408 ( 1 9 2 5 ) . Hendricks and Albrecht: Ber., 61, B, 2153-61 ( 1 9 2 8 ) . Welo and Baudisch: J. Biol. Chem., 6 5 , 215-27 (1925). In the present paper the term “hydrated Fer04”should be used to describe the precipitate formedbyadding an iron nitrate solution containing a Z : I mol ratio of Fe+++/Fe++to hot NaOH, or strong SH,OH.

’

42

P. H. EMMETT AXD KATHARINE 5. LOVE

ture of 3 0 0 T or less. The hydrated Fe304was initially prepared by adding the appropriate mixture of ferrous and ferric iron solutions to hot NaOH, and filtering, washing and drying the precipitate. The rhombohedral Fez03 was prepared by heating the cubic form in air to a temperature of 550°C or more. The relative catalytic and surface characteristics of the two oxides prepared in the above manner were then determined by ( I ) the catalytic oxidation of benzidine with hydrogen peroxide; ( 2 ) the growth of Bacterium Zepisepticum,(3) the absorption of oxygen; and (4) the absorption of water vapor a t 25’C from air saturated at the same temperature. The first three of the above comparisons revealed such a marked difference in activity between the cubic sample and the rhombohedral sample as to cause Welo and Baudisch to designate them as “active” and “inactive” Fe203. The fact that the water sorption values were approximately the same for both oxides, being in each case some 2 0 to 2 7 % of the weight of the samples used, was interpreted as sufficient evidence that the total overall surface of the two oxide samples was the same and that the differences noted above were t o be attributed to a difference in crystal structure. Accordingly, they concluded that the catalytic activity of Fe203depends on its crystal structure and may be vanishingly small for the rhombohedral modification. A priori, one might expect just such results as obtained by these authors. It has been pointed out1 that in order to convert Fe304to an isomorphous Fe203form it is necessary to crowd 4 atoms of oxygen into each unit of structure of the Fe304. Futhermore, attention has been called1f2to the fact that either of two possible positions assigned by symmetry considerations to these atoms apparently necessitates their being crowded into positions not sufficiently large to admit oxygen atoms of normal diameters without distorting the lattice. These extra oxygen atoms might be expected to be removed from the cubic Fe203lattice by reduction more easily than the normal oxygen atoms from the stable Fe203. By the same line of reasoning a marked difference in catalytic activity even to the point of one being “active” and the other “inactive” might not be surprising. However, closer examination of the procedure employed by Welo and Baudisch led the authors to the conclusion that the evidence for cubic Fe203 being “active” and rhombohedral F203being “inactive” is not conclusive. Their catalytic and surface work is open to objection due to the presence of the following disturbing factors: (I) The fact that the cubic sample of Fe203 as tested, had been heated to only 300’C. whereas the rhombohedral sample had been hented to 550’C. or higher, renders the drawing of valid conclusions as to the dependence of catalytic activity upon the crystal structure of the two materials impossible. ( 2 ) The catalytic tests used are at best qualitative and are not sufficiently diversified to warrant drawing general conclusions relative to catalytic activity. ( 3 ) The presence of impurities such as NaOH or Na2S04might materially affect the tests made. (4) The values obtained Hendricke and Albrecht: Ber., 61B, 2153-61(1928).

* Welo and Baudisch: Phil. Mag., 50, 399-408 (1925:.

CUBIC AND RHOMBOHEDRAL FERRIC OXIDE

N "

-em c c

"0 "0 O O F , 10 10

1010010

Lad

8

0 0 0

Z R R

N "

8 h

43

44

P. H. EMMETT AND KATHARINE S. LOVE

for water sorption are neither of the right order of magnitude nor takenin such a manner as to disclose the relative surface areas of the oxides. The picking up of 2 5 % by weight of water vapor by the oxides was t o a considerable extent doubtlessly hydrate formation, capillary condensation, or absorption of H t 0 by hygroscopic impurities in the test samples. This latter source of error was especially apparent in one sample that was reported to have picked up water equivalent to 95% of its own weight. Part of the water was admittedly "free and flowed about in the boat in which the dry oxide had been placed." The present research has, therefore, been carried out and consists in a study of: (I) the catalytic activity of the two forms of iron oxide prepared under identical conditions; ( 2 ) the sorption of water vapor upon each; (3) the effect of A1203 upon the rate of change of the magnetic permeability of the cubic Fe2O8,a t temperatures ranging from 4 j o to 6oo0C; (4) The rate of reduction by hydrogen of the cubic and rhombohedral samples; ( 5 ) the influence of A1203 upon the rate of reduction of cubic Fe?03to Fe304and to Fe.

Preparation of Oxide Samples To avoid several of the above-mentioned disturbing factors in the work of Welo and Baudisch, it was necessary to prepare samples both of cubic and of rhombohedral Fe203by methods of precipitation, filtration and drying that were as nearly analogous as possible. This was accomplished as described below by precipitating Fe(OH)3, drying the precipitate a t 3ooOC in air, to form the rhombohedral Fe203, and comparing it with a sample of cubic Fe20a obtained from similar oxidation of hydrated Fe304 in air a t 3ooOC. To preclude the possibility of traces of NaOH or N a ~ S 0 4disturbing the tests, samples were prepared using N H 4 0 H as precipitating agent instead of NaOH. I n addition, of course, samples were used that were prepared and treated in the same manner as those of Welo and Baudisch. TABLE I1 Analysis of the Oxides Catalyst Crystal Number Form'

Rhombohedral Cubic 3 Rhombohedral 4 Cubic 5 Rhombohedral 6 Cubic I

2

Per cent Alto3

Per cent NaOH

-

-

I . 58

1.55 -

I

Per cent Ferrous Fe

Per cent Fe(NOs)s

Per cent Total Fe as Fe103 in Sample As After ignition taken to 8 ~ 1 0 0 0 " from for 2 hrs. in bottle elec. furnace

-

"one

99.36

Xone

99.63

-

-

-

-

None

Sone None None

I2

0.12

o 07

-

-

1oo.9(1) (1)

100.7

-

Possibly the formation of some ferrous iron by ignition of the oxide is the cause of this apparent high analytical result. * The crystalline forms of the samples were determined by X-ray photographs very kindly taken by Dr. S. B. Hendricks of this Laboratory.


45

Sine samples of catalysts were prepared as indicated in Table I. I n all cases the precipitate was washed by decantation until the decanted liquid was not alkaline. Two or three additional washings were then made. A11 samples were dried in air that had passed through soda lime and finally PzOs tubes. The temperature of the drying furnace was automatically held a t 300' i IoOC. The samples were sealed in glass-stoppered bottles till ready for use. Samples of the various catalysts were analyzed for possible impurities. Samples I and z were analyzed for total iron, the original weight being that of the sample as taken directly from the sample bottle. They were then heated for z hours a t 800'-1000'C., cooled in a desiccator] and again analyzed for total Fe203. The results are shown in Table 11.

Catalytic Oxidation of Benzidine and Guajac Resin by Hydrogen Peroxide As a measure of catalytic activity, 4 reactions were made use of: the oxidation of benzidine by hydrogen peroxide in the presence of the catalyst, the oxidation of guajac resin by hydrogen peroxide in the presence of the catalyst, the catalytic combination of hydrogen and oxygen at a temperature of z50°C, and the catalytic decomposition of ozone at temperatures ranging from - 75'C to z j o c . TABLE I11 Activities of the Catalysts toward Organic Oxidation Reactions Benzidine oxidation Guajac resin oxidation

Good (a)

Fair

5J'13,216,4, 5,6, I ,3,4,2

6a 6a

Very Poor

sa1 2 % P

I

2a

(a) The "good" catalysts are arranged in order of their decreaaing activity. However, the activities of catal sts 1-6 toward both benzidine and guajac resin were so nearly the same that it waa difficd to distinguish between them. Accordingly for these oxidation reactions one is justified in concluding merely that the rhornbohedrai modification of FelOa is at least as active catalytically as is the cubic form.

The results obtained by repeating the tests mentioned by Welo and Baudisch were, due t o the nature of the tests, necessarily qualitative. They show clearly, however, as can be seen in Table 111, that the activity of the rhombohedral catalysts dried in air a t 30ooCJ (catalysts I , 3 and 5 ) is as great as, or greater than that of the corresponding cubic FezOaprepared in air &t 300' (catalysts 2, 4 and 6). Furthermore, heating either cubic Fe?O3 or rhombohedral Fez03 t o 550' destroys its activity (catalysts 2a and 5a, respectively). l Cnfortunately sample 6a as treated throughout this paper was prepared and used before noting that the several hours treatment at 5jo" had but slightly impaired its magnetic permeability. Either traces of impurities or some characteristics of the method of precipitation render this oxide less readily convertible to the rhombohedral form and make difficult and slow also the change of catalytic and, as we shall later see, other properties. However, the behavior of catalysts I , z,3 4, j, za and j a pointed out above leaves no doubt as to the catalytic activity in these readdons being dependent upon heat treatment of the oxide and not upon their crystal structure.


46

It is believed, therefore, that the difference in activity of cubic and rhombohedral Fez03 previously ascribed to crystal structure can in reality be explained by differences in the thermal treatment of the two oxides. Accordingly, in so far as these two reactions are indicative of catalytic activity, the conclusion seems warranted that both cubic and rhombohedral Fez03 have approximately the same activity. A search of the literature revealed very little information relative to the exact nature of the two catalytic reactions here cited. Hence several observations made during the present work may be of interest. It was definitely ascertained that a very small amount of ferric iron was sufficient to give the blue coloration produced in either of these reactions. I n fact, the resin test for ferric iron is of the same order of sensitivity as the well-known KSCN test. To ascertain whether traces of ferric iron from some Fe(N03)3occluded as an impurity in the oxide might not be responsible for the so-called catalytic activity manifested by samples I to 6 mentioned in the above table, thorough digestions of the oxides in water were carried out, both with and without hydrogen peroxide. Subsequent test of the filtrate from this oxide water mixture failed to give a measurable trace of ferric iron with KSCN. Similarly, tests of the filtrate with the resin and hydrogen peroxide resulted in no blue coloration. However, the addition of a small amount of catalyst t o this mixture of filtrate, H202 and resin extract, immmediately produced a deep blue coloration. Thus, it seems that the two organic oxidations are catalyzed or caused to occur either by the iron oxide itself or by some impurity not removed from the oxide by repeated washing but dispelled by heating to jsooc.

Catalytic Combination of Hydrogen and Oxygen Determination of the efficiency of the catalysts toward the hydrogenoxygen reaction was made by passing a mixture of 97.5% oxygen and 2 . j% hydrogen a t a rate of flow of 116 cc. per min. over 4 cc. of catalyst. The apparatus used is essentially that originally described by Pease and Taylor.' The oxygen was purified by passage over hot platinized asbestos, through concentrated H2S04, soda lime, and finally P2Oi. The hydrogen was generated electrolytically from a NaOH solution using platinum electrodes. The hydrogen-oxygen mixture from the cell joined the stream of oxygen gas before entering the sulfuric acid wash tube. The only rubber connection in the line was a short one between the tank of oxygen and the entrance to the purification line. The apparatus throughout was of Pyrex. The water vapor formed by the catalytic combination of the hydrogen and oxygen was determined by dehydrite-filled weighing tubes. The exit gas from the weighing tubes was analyzed for hydrogen. Great care was taken during each run to embed the thermometer the same distance in the bed of catalyst. The catalyst samples were broken to about 20 mesh size and runs a t 2 j o o were carried out. Runs at 315'C were also made, but the conversions were so Pease and Taylor: J. Am. Chem. SOC., 43, 2181 (1921)


47

high as to be of little use in differentiating between the activities of the catalysts. The results obtained in this series of catalytic runs are summarized in the tables below:

TABLEIF7 H2- O2 Catalysis Catalyst number

I 2

2a 3 4

5 .;a 6 6a

Crystal Form

Weight of Sample (grams)

rhomb. cubic rhomb. rhomb. cubic rhomb. rhomb. cubic cubic

j.16 3.26 3.43 5.24 4.I7 5.27 5.54 4.51 4.33

Apparent A Volume H,O formed of (mg./5 min.) Sample (cc) 250°C

4 4 4 4 4 4 4 4 4

B Efficiency Exit H2 in A terms of equi- = a+B valent H 2 0 / 5 min. 250'C 2jo"C

5.3 4.3 4.4

3.6 5.4

5.1

3.9 6.1

3.7 3.0 1.5

2.8 I .6

5.1

-

8.2 6.3 8.2

0.60 0.44 0.46 0.56 0.38 0.30 0.16 0.31 0.17

.I = mg. H20 formed per j min. at steady state by passage of the hydrogen-oxygen mixture U

=

over the FeZOa. mg. H 2 0 equivalent to the hydrogen not cataIyzed to water by passage over Fe?03.

It will be noted that the activity of the rhombohedral FesOais as high or higher in every case than that of the corresponding cubic Fe203as evidenced by the higher conversions obtained on passing the gaseous mixture over equal volumes of the two catalysts. However, as will be noted in Table IV, the heating of sample z to 550°C seemed to have no effect upon its ability t o catalyze the combination of the hydrogen-oxygen mixture. It is thus apparent that for some reason this reaction will not distinguish between the catalytic activities of oxides that in the organic oxidations manifest manifold differences. The probable significance of the results here obtained in the hydrogenoxygen reactions will be discussed again in connection with the water sorption iralues. Ozone Decomposition For comparing the catalytic activity of cubic Fez03 with that of rhombohedral Fe203for the decomposition of ozone, catalysts I ] 2 and za were chosen. The experiments were made.in a flow system over a temperature range of -74'to 25OC. The oxygen, purified as described above, was passed at I O O cc. per min. through an ozonizer of the type used and described by Iiarrer and W ~ l f . ~ With 85 volts on the primary of a I kw., I I O - Z ~ , O O O V. transformer] an ozone volume concentration of approximately z q c could be produced in the flowing oxygen. Iiarrer and Kulf: J. Am. Chem. SOC., 44, 2392 ( 1 9 2 2 ) .

48


The experiments were carried out in an apparatus of the type shown in Fig. I . The catalyst tubes were joined t o the main ozone supply b y paraffined joints in the way suggested by Ray and Anderegg.' Thoroughly washed and dried asbestos wads were found not to catalyze the decomposition a t room temperature and were therefore used to support the catalyst sample. Very small samples of catalyst had to be used in order to avoid complete decomposition of the ozone even in the runs a t -74OC. The exit ozone from the catalyst tube was determined in the usual way by passing it through To transfirmer

S - Asbestos plug

A - Coofinq bath B - K I mlufion

C snd,E- Glass caps,

-

paraflied on

D Glass and paraffin connection

FIG.I

potassium iodide solution, acidifying the resulting solution with H2SOa,and titrating the liberated iodine with sodium thiosulfate. The titration with . I N sodium thiosulfate solution was sensitive to *0.03 cc. of ozone. Five minute runs with a flow of I O O cc. of oxygen per minute were used throughout. The low temperatures were obtained by the hand regulation of a COZalcohol bath and were constant within a degree or so during any set of runs. The results of these catalytic runs are summarized in the Table V. The results show that catalysts I and 2 , the rhombohedral and cubic samples of Fe203,manifest practically no difference in catalytic activity a8 regards the decomposition of ozone under the conditions mentioned above. The values for catalyst za at - 740CI recorded in Table VI, indicate a rate of ozone decomposition about 65% smaller than for catalyst z Similar samples of za used in later runs were from 30 to 7 5 7 0 less active than catalyst 2. While the results seem to indicate the superiority of catalysts I and z over catalyst Ray and Bnderegg: J. Am. Chem. SOC., 43, 969 ( 1 9 2 1 ) .

CUBIC AND RHOYBOHEDRAL FERRIC OXIDE

49

TABLE V Catalytic Decomposition of Ozone The relative activities of .047 gram (.03 j cc., calculated) of Catalyst KO. I ,031 gram (.038 cc., calculated) of Catalyst KO.2 ,030 gram (.035 cc., calculated) of Catalyst KO,2a Gas flow to ozonizer = IOO cc. 0 2 per minute Gas leaving ozonizer contained approximately 2 % 03. Catalyst number

Temper$ure C

No. of consecutive 5 minute runs

I

21

6

I

- 30 - 73

5

2a

27

I2

2

26

5

2

-30 - 30 23 - 74 - 73

5

I

za

2a

2a 2

I

-t 4

2

20

5

5 9

8 6 3 6

Range of per cent Per cent Decomposition at Decomposition Steady State 100-99.9 99.9 99.2-99 .o 92.7-94.2 91.1-85.4

99 .o 94 ' 2 85.8

99.3-99.8 98.6-98.2

99.8 98.4 13.3

76.3-73.3 91.5-87.8

55

.

.2-5 I j

88.5 52.1

92.3-90.7 89.8-89 . o

90.7 89.0

99.2-98.5

98.7

Summary of the Steady State Values 23-25'c

- 3ooC

- 74OC

99.9

99.0

94.2 89.0

99.8 98 7

98.4

90.7

85.8 88.5

73.3

52; I

'

2a

2 8 , they do not indicate a difference of even the same order of magnitude as the one shown to exist in the case of guajac resin or benzidine oxidation with hydrogen peroxide. As a means of obtaining an approximate estimate of the small temperature coefficient indicated by the above experiments, a few determinations of the variation of the rate of catalytic ozone decomposition with variations in the temperatures and the ozone concentrations were made. The percent ozone in the entering oxygen stream was varied by changing the voltage on the primary of the transformer. The results, obtained upon varying the ozone content of the entering gas, are shown by the data in Table VI. Values for the apparent reaction velocity constants, obtained both on the assumption

P . H. EMMETT AND KATHARINE S. LOVE

50

TABLE T'I Order of Reaction-Catalytic Decomposition of Ozone .032 gram (.os7 cc., calculated) of catalyst KO.2a Gas flow to ozonizer = 100 cc. O2per minute Temperature = - 74OC. Per cent On initial gas stream

Per cent Os in exit gas

0.577 0.582 0.625 0.610 0.622 0.318 0.325 0,334 0.340 0.356 0.726 0.746 0,759 0.762 0.403 0.410 0.414

First Order decomposition k, .!lnEo tc P 40 ' 7

40.4 38 . o 38.8 38.2 34.8 34.1 33.2 32.6 31.1 33 ' I 32.2 31.7 31.5 27.1

26.5 26.2

Second Order decomposition k? =

f (L - !-) to

pr

PO

5.32 5.25 4.74 4.91 4.77 8.88 8.59 8.23 8.01 7.44 3.78 3.52 3.52

3 .so 6.02 5.84 5.73

that the catalytic decomposition is first order with respect to the partial pressure of ozone, and on the assumption that it is second order, are included in this table. The calculations were made by the method ordinarily used for catalytic flow systems. Thus

where V = total gas flow in cc. (S.T.P.) per min. V, = Apparent volume of the catalyst as measured approximately in a small calibrated glass tube. T = Absolute temperature of catalyst chamber. P = Total pressure in catalyst tube = I atmosphere in present experiments. pf = Partial pressure of ozone in exit gas po = Partial pressure of ozone in entering gas kl = Apparent reaction velocity constant for the catalyst (time in seconds).

CCBIC A S D RHOMBOHEDRAL FERRIC OXIDE

51

Similarly, for the second order calculations, kz = I / t c (I/pr - I/PA (2) where pf and po are expressed in mm. - -V T P -

6ovc 2 7 3

(i- i)

In both equations, t, = “time of contact,” and hence may be expressed in terms of the gas flow and volume of catalyst as shown in the equations. The above equations are identical with those that one obtains from calculations of the Langmuir type,’ assuming the reacting gas to be slightly adsorbed and the decompositions to be monomolecular and bimolecular, respectively, on the surface of the oxide. I n Fig. 2 are shown the kl and k2 values obtained from the data of Table VI. The gradual drift in the successive readings presumably is due to a slight poisoning of the catalyst a t -74OC. It does not, however, invalidate the conclusion that the decomposition follows approximately a first order equation. The agreement with the second order equation is poor. Furthermore, a moment’s reflection upon the data will show that it does not correspond to a zero order reaction. Hence, these preliminary runs indicate that the decomposition of ozone over Fe203is an apparent first order reaction. Data for additional experiments at three different temperatures are shown for catalyst za in Table VII. The values of kl for these two runs are plotted against I/Tvalues in Fig. 3 . From the slopes of these two curves and from data for several other runs an “apparent energy of activation” of 2000 + 500 calories per mol of O3 is indicated. The reaction velocity constants kl and k2are of course to be distinguished from specific reaction velocity constants as ordinarily defined for homogeneous reactions. For a catalytic apparent monomolecular decomposition uninhibited by the products of reaction one can write for the reaction at constant volume d ! = dt

where PA

t x

=

x.s

partial pressure of reactant at any time

= time in seconds = reaction rate per sq. cm. of adsorbed reactant on the surface

= Total surface of the catalyst fraction of the surface covered by reactant. The Langmuir adsorption isotherm equation, dp*- kpA dt I kp., dB= kpr reduces to the simple form dt when the adsorption of gtw A is slight. Hence

S

c

+

-dpA = XkSpA = klpr dt Similar consideratione are involved in determining kl. This equation when integrated becomes ( I ) above.,, Furthermore, application of the above equation to a flow system involves “time of contact in a manner that is in error by several h u n e d percent. Thus in place of VC in the above equations one should more correctly use V c , where V’o ejluals Vo - Vs, and VSis the actual space occupied by the catalyst particles themselves. V is usually between 25 and 7.5% as large as VC. None of these factors, however, affects the validity of the conclusions drawn above as to the apparent order of reaction and “apparent temperature coefficient” of the catalytic reaction.

P. H. EMMETT AND KATHARINE S. L O V E

Tme

In rn;nufes

FIG.2

Y T

FIG.3 Plot of r/T vB. the logarithm of the apparent reaction velocity constant for catalytic decomposition of ozone by Fe?Os.


53

TABLE VI1 Temperature Coefficient of Catalytic Decomposition of Ozone ,032 gram (.037 cc., calculated) of Catalyst No. 2a Gas flow to ozonizer = IOO cc. Gas leaving ozonizer contained Run Temperature

"C 23 24 23 23 23

- 30 - 30 - 30 - 30 - 30

0 2

2%

per minute 0 3

Run

I

Per cent Decomposition (a)

99.2 99.2 99.4 99.2 99 .o 96.8 96.0 95.4 95.5

95.3

Temperature

"C 22 22 22

22 22

-31 -31

-3' - 30 -30 -31

-30 -30

- 72

-

- 74

72.3 72.6 72.4 71.7 7 1 .o

2

Per cent Decomposition 98.o

- 74 - 74 - 74 - 74 - 74 - 74

98.2 98.6 98.2 98.1 93.1 92.8 92.2 92.7 91.2 91.3 91.4 91.6

70.7 71.2

71.8 71.8 71.8 70.9

(a) Readings in Run No. I and Run No. 2 were taken chronologically in the order listed and for a 5 minute period. Between runs No. I and No. 2 the catalyst was closed over night in dry air.

It is hoped in the near future to study more closely the kinetics of the catalytic decomposition of ozone with larger amounts of less active catalyst and a t room temperature] rather than -74'C.

H20Sorption by the Various Catalyst Samples The water sorption values were determined both a t 25OC and a t 210°C. The values were obtained by passing a stream of nitrogen saturated with water vapor a t OOC., and containing accordingly 0.5770 water vapor, over the catalysts dried at 257OC. in situ immediately before each run. The water vapor in the exit stream of nitrogen was determined. Passage of the gas was continued until the exit and the entering water vapor contents were the same. The difference between the water content of the influent and effluent nitrogen stream during this saturating period was taken as the

54

P. H. EMMETT A S D KATHARINE S. LOVE

. .

OIb v) roffi

h

m. v .)

. .

CAN CI

O N ^

F:-

b

-

0. -.

w< N

?t

N W

h

* H

3

. . N

. N. *. -.

vi

w e

Ne

hm

ev,rovi

--

m.

O

N

. .

10

00

N

- -

eo1

. .

"9 N

N

4'9 N

H

3

"? -e

h

"09

"

w vie "1

k

i

O d O c N

m

ee O e N

OIC.

. .

. .

O m

h

m. - . *OI

e

N N

. . *.

vi-

v)W

v)

h O I h h

. . . .

w e e e N

tt

N N

mr?

Oe

vi vi

309 0909

""9 i

w

h

? ? i h h

mm

N N

h h

. w10 . . e. 509

vi

KJ2

vi0

w o

N

N IN 3

v)

3

W

09 ? "

o wm ")

. . '9'9'0

r?n:

"1 c,

N

*

N e vi

w, w

ee*e

mrc:

N

. . .

N N v i

rC,r?r?rc)

d b

z

N

N

N

Virnro

090909 N "


55

“sorbed” water. This value was in every case checked by removing the water a t 257OC. until the exit water was less than I mg. per hour. The values showed very good agreement with the check in every case. The results obtained for the different catalyst samples are shown beloq in Table T’III. From the water sorption values the following conclusions are apparently justified : (I)

are

IO

The water sorption values per gram of catalyst for cubic Fe,Os samples to 6 0 7 ~smaller than those for the rhombohedral form.

(2) The sorption on samples za and sa is about 1/4as great as that on the samples not heated a t 5so°C. Sample 6 withstands heating so well that it is not materially changed by a two hour exposure to a temperature of sso°C. Hence its water sorption value has changed but little.

(3) Previous use of sample za or sa for catalysis of a hydrogen-oxygen mixture does not increase the water sorption values.

It should be noted that whereas the activity of catalyst za as gauged by the hydrogen-oxygen catalysis is as high as that of catalyst z , the water sorption value is only about 1 / 4 as large. This indicates the possibility of a sintering process being able to close up pores and crevices that play an important part in the “sorption” of water vapor, but not in the catalysis of a hydrogen-oxygen mixture. Magnetic Permeability Determinations The rate of change with heating of magnetic permeability of samples 2 , 4, and 6 was determined by a method identical with that outlined by Welo and Baudisch. The experiments were made by alternately measuring the permeability of a particular oxide and heat treating it for an hour at some definite temperature. Such hourly heating periods were continued until the magnetic permeability became 1.05 or less. The initial permeabilities of catalyst 2) 4 and 6 were 1.93, 2 . 2 8 and 2.48, as measured in a magnetic field intensity of zoo Gauss. These permeability measurements agree qualitatively with those of Welo and Baudisch. The results of the measurements are summarized in Fig. 4. From the curves it is evident that catalyst 4 is definitely much slower in changing from the cubic to rhombohedral form, or, more correctly, in changing from a permeability of 2.3 to 1.05, than is catalyst 2. Both of these materials were prepared by identical methods, except for the addition of the promotor to KO.4; it would seem, therefore, reasonable to ascribe the inhibition of the rate of transformation to the promoter content. The rate of change of the magnetic permeability of catalyst 6, the Lefort’s oxide, made by precipitation with NaOH, is very much slower than that of either sample precipitated yith “,OH. I t will be noted in connection with Table I that catalyst 6, in spite of careful washing, was found to contain about . I Z % NaOH.

56

P. H. EMMETT AND JCATXMINE S. LOVE

Effects produced by this impurity, or by a change in particle size with the precipitating agent, seem to be. the most likely causes af the differtmaebetween the rate of c h g e of magnetic perm&bilityof catalyst 6, a d that ,of either catalyst 2 or 4

FIG.4 The influence of promoter content and recipitating agent on the rate of decrease of permeahlity of cubic FegO, in &e temperature range 4~0°-600nC.

Reduction of Cubic and Rhmbcihdral l?efi3 The comparative rates of reduction of the oxidees, in a stream of 2 0 0 cc. of hydrogen per minute, were determined by a flow system &dar fo t b t described by Pease and Taylor. Hydrogen was pwiiied b y passage over hot copper and dried by passing through soda lime and P& t u b s . The sample of oxide immediately before the reduction was dried t o constant weight at z70°C. Water was determined in the exit gases by colleetiqit in P206tubes. The rate of reduction is expressed in terms of mgs. of water obtained per 5 minutes. The temperature of the sample was obtained from a calibrated thermometer thrust thrciugh the top of the furnace into the bed of catalyst. Typical sets of reduction curves are shown in Fig. 5 and Fig. 6. From Fig. Sa, gb, and gc, it is evident that the reduetiion xate ~ u w e sfor r k t m b o M a l and cubic Fez03 are qualitatively the same. Unfortunately, the difference in size of the two samples used makes a quantitative comparison between the two rates difficult. However, if due allowance be made for the &e of the sample, one is led t o the conclusion that the cubic modification nduces a little more rapidly than the rhombohedral. The same figures also show that the rate of reduction of the samples that had been heated t o 55oOC is very much smaller than that of those dried at 300°C. Thus samples za and s a each reduces at a very much slower rate than do samples z and 5 , respeetively. Sample 6a also reduced a t a somewhat lower rate than sample 6. However, since the few hours heating at 550' had but little effect vn the magnetic

CUBIC AND RHOMBOHEDRAL FERRIC OXlDE

57

T 6 7%

9 4 3 2

1

4:

3

I

4

5

6

T/me in Hours

FIG.ja Rate of reduction by hydrogen of cubic and of rhombohedra1 FezOa at

210°C.

,

FIG.5b Rate of reduction by hydrogen of promoted cubic and rhombohdedral FezOl at ZIOOC.

0 - E x p 244-Cdi”o.5{ 5.42 grams 4. I cc fCd/C) grams x-€xp 22- Cat No. 6 4.04 3 . 6 cc (ralr)

{

F 4

2

FIG.5c Rate of reduction by hydrogen a t 210°C of cubic and rhombohedral samples of Fe?03in the preparation of which S a O H was used.

58


permeability, crystal structure, or other characteristics of this catalyst, it was to be exptcted that its reduction rate would be affected less markedly than in the case of sample z or 5 .

T,me in hours

FIQ.6b The influence of N1O3on the reduction by hydrogen of FesOl to Fe.

Several reductions of cubic Fe203 with and without the 1.576 A1203 were run, first at 24ooC., and finally to completion a t 345OC. The results of these runs are illustrated by typical examples in Figs. 6a and 6b. The reduction of Fe203to Fe304is changed but little, as shown in Fig. 6a, by the presence


59 \

of the A l 2 0 3 . On the other hand, as shown in Fig. 6b, the reduction of Fe304to Fe is very much slower for catalyst 4, containing the A1203, than for No. 2, the pure iron oxide catalyst. In this same figure, the curves for experiments 3 and 4 contrast therate of reduction a t z9ooC. of FetOl to Fe without and with A1203 present. Only the FesOl to Fe stage of reduction is shown for these two curves. I n all of the 210’ iron oxide reduction curves it will be noted that the first few points at the beginning of each curve are lower than later ones. This is apparently caused by the rapid sorption by the dry oxide of water formed by reduction. The magnitude and rate of water sorption values on these oxides are such as to confirm this view.

Discussion of Results It is difficult to offer an explanation or mechanism for the very marked decrease in catalytic activity toward the oxidation of benzidine or guajac resin by hydrogen peroxide produced by heating either cubic or rhombohedral Fe203samples to 550’ or higher. The phenomenon can perhaps best be explained by merely postulating the sintering of points or regions of the catalyst surface usually catalytically active toward these reactions. It is difficult to exclude the possibility, however, that the effective catalytic agent in these two reactions might be some trace of inorganic salt, the quantity of which is too small to be detected analytically, and the effect of which is destroyed by heating to 550’ or more. Exposing catalyst z to a temperature of 55oOC. for two hours (thereby converting it into what has been designated as catalyst za) produced interesting and unexpected results. The water sorption capacity of the sample was reduced by 50 to 75%, the rate of reduction by hydrogen by 7 5 to go%, and yet the catalytic activity toward a mixture of hydrogen and oxygen remained unimpaired. Of course, i t is well known that in general the surface of a catalyst as measured by the adsorption of a gas does not necessarily give information as t o the extent of active catalytic surface. Furthermore, many cases are known in which heat treatment, or poisoning, of a catalyst reduces the catalytic activity markedly without materially affecting the capacity of the catalyst for adsorbing gases. However, the occurrence of the converse phenomenon, the decrease of “sorption” capacity of a catalyst without an attendant decrease in catalytic activity is very unusual. The decrease in the reduction rate of the oxide without a change of catalytic activity is similarly unexpected. The experimental facts are in agreement with the following explanation: Each particle of the oxide (one mm. or so on a side, depending merely on the size of the screens used in preparing the original oxide) contains a certain number of pores and capillaries through which a free flow of gas does not take place. Such capillaries, therefore, would constitute a negligibly small part of the catalytic surface. However, both in the sorption of water and in the reduction by hydrogen, the capillaries would play an important

60

P. H. EMMETT AND KATHARINE €3. LOVE

rble. Accordingly, heat treatment sufficiently intense to decrease the number and size of the capillaries would a t the same time markedly affect. both the water sorption by and rate of reduction of the oxide. Though this explanation of the observed facts may not be the correct one, it is a t least plausible and consistent with all data obtained. It is of interest to compare the results obtained in the present work on the catalytic decomposition of ozone by Fez03with some observations made by Strutt.1 He noted the rate of decomposition of ozone a t low pressure over silver oxide and calculated that it corresponded to the decomposition of every molecule of ozone striking the silver oxide. If this calculation be correct, it follows that the temperature coefficient of the decomposition would be small. Since there seems to be no good reason for expecting the catalytic decomposition by silver oxide to differ markedly from that by ferric oxide, the low temperature coefficient and extremely fast reaction rate upon Fez03is not unexpected. The temperature coefficient for the bimolecular homogeneous decomposition of ozone corresponds to an energy of activation, as calculated by the Arrhenius equation, of 30,000 calories per mol of ozone according to Wulf and Tolman.2 The two thousand calorie apparent energy of activation for the catalytic decomposition is a much smaller fraction of the homogeneous value than is usual for catalytic reactions. Ordinarily, the temperature coefficient for the catalytic decomposition of a material is about one-half that of the homogeneous decomposition of the same reactant. The part played by the Al208 in catalyst 4, in decreasing the rate of change of magnetic permeability of a sample of cubic Fez03 is not entirely clear. It seems quite probable that the phenomenon is similar to the retarding effect of A1203 on‘the rate of growth of iron crystals in catalysts used for synthetic ammonia production3 or to the retarding effect of A1203 upon the rate of decrease of pyrophoricity of iron catalysts when the latter are heated a t temperatures in the neighborhood of 60ooC. as observed by Tammann4 and his coworkers. I n conclusion, it is well to note one possible inference to be drawn from the oxide reduction work. Benton and Emmetts in studying the rates of reduction of NiO and of Fez03by hydrogen, found the former to be autocatalytic and the latter to be “non-autocatalytic.” This result was interpreted as indicating the existence of solid solutions between Fez03and FeaOain accordance with Langmuir’s6 generalization that in such a reduction, the “active mass” of an oxide will be its total surface in case the two solid phases form solid solutions, but will be the interface between the two solid phasev in case the solids form separate phases. However, x-ray crystal structure work would lead

* Strutt: Proc. Roy.

SOC.,87A, 302 (1912). Rulf and Tolman: J. Am. Chem. SOC.,49, 1650-64 (1927). Wyckoff and Crittenden: J. Am. Chem. SOC.,47, 2866-76 (1925). Tammann and Nikitin: J. RUM. Phys. Chem. SOC.,56, 115-9;Sikitin: Benton and Emmett: J. Am. Chem. SOC.,46, 2728-37 (1924). Langmuir: J. Am. Chem. SOC.,38, 2263 (1916).

121-7 (1925).


61

us to believe that the formation of a series of solid solutions between rhombohedral Fe203and cubic magnetite is highly improbable. Accordingly, the present reduction work in which cubic Fe20sis reduced, becomes of particular interest, since this latter material should be capable of forming a complete series of solid solutions with magnetite. With such a state of affairs, it is pertinent to inquire whether the reductions of rhombohedral and cubic Fe203 exhibit any marked differences in the nature of the reduction curve. As can be seen in Figs. sa, gb and gc, no appreciable difference in the nature of the two oxides is found to exist. This it seems can be interpreted as indicating either (I) that the removal of oxygen from the oxide is not the rate-determining step in the reduction, or ( 2 ) that Fea04forms solid solutions as rapidly and easily with rhombohedral Fe203as with cubic Fe203, or ( 3 ) that contrary to the inference usually drawn from Langmuir's hypothesis, reductions or heterogeneous reactions can be autocatalytic or non-autocatalytic regardless of whether solid solutions or separate solid phases are formed. summarp

-4 comparison of some of the chemical and catalytic characteristics of cubic Fez03with those of rhombohedral Fez01hasrevealedthefollowingfacts: ( I ) The catalytic activity of each of the two similarly prepared oxides, toward the oxidation of either benzidine or of guajac resin by hydrogen peroxide is approximately the same and is not, accordingly, a function of the crystal structure of the two oxides. Heating either cubic or rhombohedral Fe203 to 5go0C. for several hours practically eliminates their catalytic activity toward these oxidations. (2) The catalytic activities of both rhombohedral and cubic Fe203are approximately the same a t 250°C as measured by the rate of combination of a hydrogen-oxygen mixture on their surfaces. Heating either sample to 550' does not materially alter its catalytic activity toward this reaction. ( 3 ) The catalytic decomposition of ozone a t -74OC occurs equally rapidly on the rhombohedral and cubic forms, but less rapidly on the highly heated material than on that dried a t 300°C only. The decomposition of ozone in a 2% ozone-98% oxygen gas stream at an overaIl space velocity of about 180,000 to zoo,ooo is practically complete on the most active of the iron oxides, even a t -74°C. The reaction has a temperature coefficient corresponding to an apparent energy of activation of approximately 2 0 0 0 calories. It approximates an apparent first order decomposition. (4) The water sorption from a stream of nitrogen containing . 5 7 % water vapor has been measured at both 2 joand 210'. The results show: (a) The water sorption on a given weight of cubic Fe20s is I O to 60% less than that of the rhombohedral material. (b) The sorption values on samples of the cubic and rhombohedral Fe20a that have been heated for several hours at 550°C. are 7 j-80% less than on the samples dried at 300OC.

62


(c) The water sorption on a highly heated sample of Fez03is not increased by activation with a hydrogen-oxygen mixture. ( 5 ) The loss of magnetic permeability on heating is most rapid for a cubic Fea03made from hydrated Fe304precipitated by NH40H, next fastest on a sample similarly prepared except that it contains 1.5% Al2O3,and slowest on the cubic Fe203 made from hydrated Fe304precipitated by NaOH. (6) The rates of reduction by hydrogen of the two types of Fe203are about the same, being a little faster for the cubic than for the rhombohedral forms. The rates of reduction are decreased by from 50 to go% when either the cubic or rhombohedral form is heated to 550°C. for two hours. ( 7 ) A1203 precipitated with Fe304, and hence contained in the final cubic Fe203,has little effect on the rate of reduction of Fe203to FesOa, but seems to markedly retard the rate of reduction of Fe304to Fe. I n conclusion, the authors wish to express their thanks to Dr. Sterling B. Hendricks, for his kindness in taking the necessary x-ray powder photographs, to Mr. C. W. Gelhaus, for having analyzed all the oxides used in the present research, and to Dr. 0. R. Wulf, for his many suggestions relative to the technique of producing and working with ozone.

A Comparison of the Catalytic and Chemical Characteristics of Cubic

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