Mechanisms of the Room Temperature Oxidation of Carbon Monoxide

48 Mechanisms of the Room Temperature Oxidation of Carbon Monoxide on Nickel Oxide

Downloaded by CORNELL UNIV on May 22, 2017 | http://pubs.acs.org Publication Date: January 1, 1968 | doi: 10.1021/ba-1968-0076.ch048

G. EL SHOBAKY, P. C. GRAVELLE, and S. J. TEICHNER Université de Lyon, Institut de Recherches sur la Catalyse, Villeurbanne, France Nickel oxide, prepared by dehydration of nickel hydroxide under vacuum at 250°C. [NiO(250)], presents a greater activity in the room-temperature oxidation of carbon monoxide than nickel oxide prepared according to the same procedure at 200°C. [NiO(200)], although the electrical properties of both oxides are identical. The reaction mechanism was investigated by a microcalorimetric technique. On NiO(200) the slowest step of the mechanism is CO+ CO + Ni --> 2 CO (g) + Ni , whereas on NiO(250) the rate-determining step is O-(ads) + CO(ads) + Ni --> CO (g) + Ni . These reaction mechanisms on NiO(200) and NiO(250), which explain the differences in catalytic activity, are correlated with local surface defects whose nature and concentration vary with the nature of the catalyst. 3(ads)

(ads)

3+

2+

2

3+

2

2+

T ^ h e oxidation of carbon monoxide on nickel oxide has often been investigated (4, 6, 8, 9, 11, 16, 17, 21, 22, 26, 27, 29, 32, 33, 36) with at tempts to correlate the changes in the apparent activation energy with the modification of the electronic structure of the catalyst. Published results are not in agreement (6,11,21,22,26,27,32,33). Some discrepancies would be caused by the different temperature ranges used (27). However, the preparation and the pretreatments of nickel oxide were, in many cases, different, and consequently the surface structure of the catalysts—i.e., their composition and the nature and concentration of surface defects— were probably different. Therefore, an explanation of the disagreement may be that the surface structure of the semiconducting catalyst ( and not only its surface or bulk electronic properties ) influences its activity. A

292

Mayo; Oxidation of Organic Compounds Advances in Chemistry; American Chemical Society: Washington, DC, 1968.


48.

E L SHOBAKY E T A L .

Carbon Monoxide on Nickel Oxide

293

The catalytic activity of a given catalyst towards a given reaction combines two aspects, chemical and physical. The chemical aspect is related to the energy spectrum of species chemisorbed on a surface and involved in the reaction. The position of the energy levels of these species is determined primarily by the chemical nature of these species and the chemical nature of the surface. For a given molecule on the catalyst surface, which is most often nonhomogeneous, the position of the energy levels depends on the nature of the adsorption centers : ions or lattice defects of thermal origin or resulting from the previous history of the catalyst. Therefore, molecules of the same kind, chemisorbed on a nonhomo geneous surface, bring forth a spectrum of local levels of various nature and energy. The physical aspect of the catalytic activity of a semiconducting oxide is related to the position of the Fermi level, which depends on the state of the system as a whole and is a collective property. The Fermi level would remain the controlling factor in catalysis only as long as the local levels of the chemisorbed species involved i n the reaction remain fixed. This means that the energy spectrum of the surface is defined and remains unchanged for a series of catalysts with different Fermi levels. The purpose of this paper is to show that i n a comparison between two nickel oxide catalysts which differ i n the energy spectrum of their surfaces, the Fermi level ceases to characterize the activity. In other words, the energy spectrum of the surface becomes the determining factor i n catalytic activity. This is because after different treatments of the N i O catalyst, which are explained i n detail, the energy of interaction with the solid of chemisorbed species varies greatly, whereas the Fermi level may or may not vary. The influence of the surface structure upon the catalytic activity is likely to be particularly important in the case of finely divided nickel oxides, prepared at a moderate temperature, which present catalytic activity for this reaction at room temperature. In a previous work, we studied the room-temperature oxidation of carbon monoxide on nickel oxide prepared by dehydration of the hydroxide under vacuum ( ρ = 10" torr) at 2 0 0 ° C , by means of a microcalorimetric technique (8, 20). The object of this work is to re-investigate, by the same method, the mecha nism of the same reaction on a nickel oxide prepared at 250°C. [NiO(250)] instead of 200°C. [ N i O ( 2 0 0 ) ] . 6

Experimental Materials. The nickel oxide was prepared by decomposition of a very pure nickel hydroxide (25) in vacuo ( ρ = 10~ torr) at 250°C. (24, 6


294

OXIDATION OF ORGANIC COMPOUNDS

II


35 ). Composition, stoichiometry, electrical conductivity, apparent energy of activation of conductivity, surface area, and color of the oxide are presented in Table I together with the values for the oxide prepared at 2 0 0 ° C , according to the same procedure. NiO(200) presents a small excess of oxygen, whereas NiO(250) contains metallic nickel (Table I ) . Magnetic measurements (15) have confirmed the chemical analysis (19). However, both oxides are p-type semiconductors, as shown by the Seebeck effect measurements. In the case of N i O ( 2 5 0 ) , this result means that, although there is a total excess of nickel, the oxide phase still contains a small excess of oxygen (13). The electrical properties of both oxides are identical. Table I.

Properties of Catalysts NiO(250)

NiO(200)

Property

Composition

NiO, 0.155 H 0 (15) 2

0.016 at O

Stoichiometry

e x c

. % NiO

(19)

Electrical conductivity in vacuo, (ohm cm.)" 24°C. 200°C.

NiO, 0.11 H 0 (15) 2

0.033 at N i % NiO (19)

1

10" 1.6 Χ 10"

13 10

10" 2 Χ 10"

13 10

Apparent energy of activation of conductivity, kcal./mole

24

24

Surface area, sq. meters /gram

142 ± 7

156 ± 7

Yellow-green

Yellow

Color

Methods. The differential heats of adsorption of reagents and the differential heat of their interaction on the nickel oxide surface were measured in a Calvet microcalorimeter with a precision of 2 kcal. per mole. The apparatus has been described (18). For each adsorption of a single gas, small doses of gas are allowed to interact with a fresh nickel oxide sample ( 100 to 200 mg. ) placed in the calorimeter cell maintained at 30 °C. A t the end of the adsorption of the last dose, the equilibrium pressure is, in all cases, 2 torr. Duplication of any adsorption experiment on a new sample gives the same results within 2 kcal. per mole of heat evolved and 0.02 cc. of gas adsorbed per gram. Electrical conductivities of the nickel oxide sample are measured in an electrical conductivity cell with platinum electrodes ( I ) by a d.c. bridge. Results and Discussion Adsorption of Reagents and Product of Reaction. C H E M I S O R P T I O N OF O X Y G E N . During the adsorption of the first doses of oxygen on N i O (250), the equilibrium pressure being of the order of 10" torr, the 3


48.


295


color of the oxide changes from yellow to black. A t the same time, the electrical conductivity of the sample increases, and at the end of the adsorption (2 torr), reaches a value of 1.8 Χ 10" (ohm cm.)" . This value does not change when the oxide, after the adsorption, is evacuated at 30°C. The same results were obtained for NiO(200) (23). Therefore, irreversible adsorption of oxygen produces ionic species on both oxides, the most probable structure for these ions being 0 " , (14, 20, 23, 37). There is also, on both oxides, a reversible adsorption of oxygen which does not influence the electrical conductivity of the sample and which can be attributed to oxygen in a molecular form (23). Therefore, there is a close similarity between the processes of oxygen adsorption on both oxides. However, calorimetric measurements show that the reactivity of both surfaces, towards oxygen, is not identical. O n N i O (250) the heat of adsorption of,oxygen decreases rapidly with coverage (Figure 1). This has not been observed for NiO(200) (20), for which the initial heat of 60 kcal per mole (Figure 1) decreases to 10 kcal. per mole only at the end of the adsorption experiments. The N i O (250) surface presents active sites since the initial heat of adsorption is 80 kcal. per mole. These sites do not exist on N i O (200). Finally, N i O (250) chemisorbs less oxygen than N i O (200). 5

1


a d s )

When nickel oxide is prepared under vacuum at 2 5 0 ° C , a small reduction of the solids occurs, giving some metallic nickel. It has been shown (35) and confirmed recently, that this reduction is not caused by grease vapors or adsorbed organic materials. The reduction is probably limited to the surface of N i O and is a consequence of the poor state of organization of the surface layers of the divided oxide, resulting from the low temperature of its preparation. The first step of the reduction is the departure of oxygen which leaves ionized vacancies in the surface lattice layers. Electrons from the vacancies are then trapped by nickel ions, and nickel atoms migrate and form nickel crystallites (13). Magnetic measurements (15) have shown that this metallic nickel (Table I) which exists in nickel oxide is not oxidized at 30°C. The active sites for the oxygen adsorption, which are found on the surface of N i O (250) but not of N i O (200), are to be identified with anionic vacancies because this high heat of adsorption is not caused by the sorption of oxygen on the nickel phase (13). The decrease in the capacity for adsorption of oxygen at 30°C. when the temperature of oxide preparation is increased from 200° to 250°C. is explained by the reduction of surface nickel ions, sites for the adsorption only of oxygen, and the for mation of nickel crystallites whose surface atoms may be active towards the adsorption of oxygen at 30°C. Recession of nickel ions below the surface for N i O (250) may also contribute to this decrease.


296


II

80

60 NiO (250)

Φ

ι ε


Ο

40

20

t t L 2

cm /g 3

60 NiO (200)

D

40

J* Ο

2

0

I

1 1 1 I

0

Figure 1.

!

J

I

I I

1

I

l—LJ 2

•

crf/g

Adsorption of oxygen on NiO(250) and NiO(200)

CHEMISORPTION OF CARBON

MONOXIDE.

Chemisorption

of

carbon

monoxide on N i O (250) does not change the electrical conductivity of the sample. The same result was obtained for NiO(200) (23). The curve of differential heats of adsorption of carbon monoxide on N i O (250) presents many similarities with the curve recorded i n the case of NiO(200) (20). However, a few differences are noted. The heat of adsorption of the first dose (0.08 cc. per gram) of carbon monoxide on NiO(200) is high (42 kcal. per mole) (Table I I ) . The adsorption of the next dose on the same oxide releases only 29 kcal. per mole. The initial high value of the heat adsorption was explained by the interaction of C O with excess surface oxygen (Table I ) , giving C0 iads) (14). In the case of N i O (250), the initial heat of adsorption amounts to 29 kcal. per mole (Table I I ) . It seems, therefore, that the surface excess oxygen 2


48.


297


which reacts with C O in the case of N i O (200) does not exist on the surface of N i O (250). This oxygen cannot be identified with lattice anions, which do not react with C O at 30 °C. to form C 0 , but with strongly adsorbed species (14). It is probable that during oxide preparation under vacuum at 2 5 0 ° C , this excess oxygen is desorbed from the oxide surface. 2

Table II.

Heats of Adsorption of Oo, CO, and C 0 on Catalysts 2

NiO(200)

NiO(250)

Oxygen Initial heat of adsorption, kcal./mole Amount adsorbed at 2 torr, cc./gram

60 2.03

80 1.9

Carbon monoxide Initial heat of adsorption, kcal./mole Amount adsorbed at 2 torr, cc./gram

42-29 4.5

29 5.5

Carbon dioxide Initial heat of adsorption, kcal./mole Amount adsorbed at 2 torr, cc./gram

46-31 9.4

29 9.7


Gas

W i t h the exception of the high initial heat of adsorption of C O on N i O (200), the differential heats of adsorption as a function of the amount of C O adsorbed are similar for both catalysts. Metallic nickel which exists in the sample prepared at 250°C. may chemisorb carbon monoxide (15). However, the metal content is small and cannot account for the heat released in these experiments on N i O (250), since the heat of chemi sorption of C O on metallic nickel is still higher (42 kcal. per mole) than the heat registered during adsorption of the first dose (29 kcal. per mole). C H E M I S O R P T I O N O F C A R B O N D I O X I D E . The curve of differential heats

of adsorption of carbon dioxide on N i O (250) is similar to the curve which was obtained for the same adsorption on N i O (200) (20). The adsorbed amounts are almost identical (Table II), and the decrease in the heat of adsorption with coverage is the same for both oxides. Carbon dioxide is adsorbed at 3 0 ° C , on both oxides, in larger amounts than oxygen or carbon monoxide. It is therefore probable that the small differences which appear in the stoichiometry and the structure of the oxides when the temperature of their preparation is different do not have much influ ence on the fairly large number of adsorption sites for C 0 which exist on both surfaces. However, the high heat, which is measured for the adsorption of the first dose of C 0 on NiO(200) (Table II) (46 kcal. per mole ) and which has been explained by the interaction between C 0 and the surface excess oxygen giving COsuds) ions (20), is not produced in the case of N i O (250). This difference, previously noted in the case of the adsorption of C O (Table II), is again explained by the departure of 2

2

2


298

OXIDATION

OF

ORGANIC

COMPOUNDS

II

surface excess oxygen when the temperature of N i O preparation is higher ( 2 5 0 ° C ) . Nickel hydroxide does not adsorb the gases used in this work. Adsorption of a gas on a partially dehydrated nickel hydroxide is reduced proportionally to the water content in comparison with the amount ad sorbed on a final sample of N i O (250). Therefore, no hydroxyl groups remaining on the surface seem to be involved in the chemisorption process. As for N i O (200), the adsorption of C 0 does not change the elec trical conductivity of the sample. O n both oxides, the adsorption of carbon dioxide is mostly irreversible. Interactions between Reagents on Oxide Surface. W h e n the catalyst is exposed to a reaction mixture, the adsorbed species formed may react between themselves and with the molecules from the gas phase. A n y combination of interactions may represent possible steps of the mecha nism of the reaction. In order to isolate elementary steps, the reagents are adsorbed in successive sequences (20, 23, 35). The amount of heat evolved for each sequence gives evidence of a specific interaction if different possible thermochemical equations are compared with the thermochemical data for the homogeneous reaction. SEQUENCE I: C O — 0 — C O . Nickel oxide, covered with carbon monoxide and placed at 30°C. under vacuum (amount of irreversibly adsorbed C O , 4.5 cc. per gram ) is exposed in the calorimeter to gaseous oxygen. The differential heat registered as a function of coverage by oxygen is shown in Figure 2A. It is much higher than on pure N i O (250) (Figure 1). This gives evidence of an interaction between C O i s ) and oxygen. During the adsorption of oxygen, the electrical conductivity of the sample increases from ~ 1 0 " to 1.6 Χ 10~ (ohm cm.)" . Therefore, ionic species are formed on the surface of the catalyst. These species cannot be identified with 0 " because of their high heat of formation [120 kcal. per mole compared with 80 kcal. per mole for the adsorption of oxygen on pure NiO(250) (Table I I ) ] . The stoichiometry of the interaction, determined by comparing the amounts of each adsorbed gas, indicates that one oxygen molecule is adsorbed for almost every preadsorbed C O molecule. Therefore, as for NiO(200) (34, 35), C08"(a ions is not explained by calorimetric data but is related to the equilibrium of Reaction 2 or 2'. The transformation of C0 ~(ads) would be increased for an increased pressure of C O (30). 3

( a d s )

3

8

The same reaction mechanism (I) has been proposed for NiO(200) (8, 20). O n this catalyst also a fraction of C0 "(i ns) does not react with carbon monoxide to form carbon dioxide. However, for N i O (200), the reason for the nonreactivity of some C0 " ads) ions has been deduced from the calorimetric measurements (8, 20). 3

3

0

(

S E Q U E N C E II: 0 — C O . Oxygen is first adsorbed on N i O (250) at 30°C. The sample is then evacuated at 30°C. (amount of irreversibly adsorbed oxygen, 1.9 cc. per gram), and carbon monoxide is adsorbed at the same temperature (Figure 3). The electrical conductivity of nickel oxide containing preadsorbed oxygen 1.8 Χ 10" (ohm c m . ) ' decreases during the adsorption of C O , and at the end of the adsorption, is identical to the conductivity of the pure oxide. Moreover, carbon dioxide is con densed in the cold trap. This shows that all ionized species are trans formed into neutral species at the end of the interaction. 2

5

1


302


II

The possibility of the intermediate formation of C(V(ads) ions and, finally, of C 0 > or C 0 may be examined by the use of the follow ing thermochemical cycles: 2(g

2 ( a d s )

Cycle 1, Formation of C 0 ~ 3

0 (g) + Ni 2

2 +

= 2 0"

( a d 8 )

+ 2 Ni

( a d s )

0 =0 3 +

+80 kcal.

0 = 0.5 Θ +23 kcal.

+72 kcal.

+60 kcal.

Μ

Hypothesis: 2 0Ni

+ CO(g) + 2 N i + Ni


C0 3

2 +

= C0 -

3 +

( a d s )

3

( a d 8 )

+

8 +

( a d s )

+ Ni

CO(g) = C O

3 +

= CO

( a d s )

+ O (g) + Ni ^ 2

a

( a d s )

- 1 2 0 kcal.

- 1 1 0 kcal.

+32 kcal.

- 2 7 kcal.

+40 kcal.

+11.5 kcal.

+72 kcal.

+60 kcal.

Cycle 2. Formation of C 0 ( g ) 2

i 0 (g) + Ni 2

= 0-

2 +

+ Ni

( a d s )

3 +

Hypothesis: 0-( ) + CO(g) + N i βββ

= C0 (g) + Ni

3 +

2

2 +

CO(g) + i 0 ( g ) = C 0 ( g ) 2

2

Cycle 3. Formation of C 0 ( 2

I 0 (g) + N i 2

= 0~

2 +

a d S

+ Ni

(ads)

+112 kcal.

+71.5 kcal.

+40 kcal.

+11.5 kcal.

+72 kcal.

+60 kcal.

- 2 9 kcal. +83 kcal.

- 2 6 kcal. +45.5 kcal.

)

3 +

Hypothesis: 0-

i a d s )

+ CO(g) + N i

3 +

= C0

2 ( a d s )

+ Ni

2 +

C0 = C0 (g) CO(g) + i O ( g ) = C 0 ( g ) 2 ( a d s )

2

a

2

When the surface coverage is low (0 = 0), the formation of C 0 ~ ) ions is possible, the heat of adsorption of carbon monoxide deduced from the cycle (32 kcal. per mole) being close to the value given by a direct experiment (29 kcal. per mole) (Table I I ) . Moreover, this is the only possibility since Cycles 2 and 3 are not balanced for 0 = 0. Cycle 2, for gaseous C 0 , is balanced when the surface coverage is about half the maximum coverage by the oxygen, and this is the only cycle to be bal anced for high surface coverages. Therefore, the formation of C 0 ( g ) , from the interaction between adsorbed oxygen and carbon monoxide, is probable for high surface coverages, and indeed C O j is found i n the cold trap. Finally, Cycle 3, for adsorbed C 0 , is balanced for neither 0 = 3

(adS

2

2

2


48.



303

0 nor θ = 0.5 θ M* However, Cycle 1 is balanced only for low coverages, when the heat of adsorption of oxygen is 75 to 80 kcal. per mole. W h e n the heat of adsorption of oxygen is smaller—i.e., for oxygen coverages 0 < θ < 0,5 θM—Cycle 3 (and no other) becomes balanced.

60

•8 Downloaded by CORNELL UNIV on May 22, 2017 | http://pubs.acs.org Publication Date: January 1, 1968 | doi: 10.1021/ba-1968-0076.ch048

Ε \ 40

S

.2

Ο

20

4

Figure 3.

cm /g 3

Adsorption of carbon monoxide on nickel oxide con taining preadsorbed oxygen

It can be concluded from the calorimetric data that the interaction between adsorbed oxygen and carbon monoxide is a multiple process. For low surface coverages—i.e., on the most active sites—C(V ions are formed. During the adsorption of the next doses of C O , these ions are converted into carbon dioxide, since the electrical conductivity of the sample finally returns to its initial value. W h e n the coverage of the surface increases, a second type of interaction between adsorbed oxygen and carbon monoxide yields C 0 d s > directly, and for a still increasing coverage—i.e., on sites of a low activity—gaseous carbon dioxide is formed. 2(a

It would then appear that a second mechanism ( II ) is also probable for the catalytic oxidation of C O on N i O (250): Mechanism II i O ( g ) + Ni-* = 0 " a

0'

( a d s )

+

( a d s )

+ Ni»*

+ CO(g) + Ni»* = C 0 ( g ) + N i * 2

The same sequence of adsorptions has been studied on N i O (200) (8, 20). The interaction between oxygen preadsorbed on N i O ( 2 0 0 ) and carbon monoxide yields only adsorbed carbon dioxide. Therefore, on N i O (200), gaseous carbon dioxide is produced during the catalytic reac tion through Mechanism I (8, 20), whereas on N i O ( 2 5 0 ) two reaction paths are probable (Mechanisms I and I I ) . These results show clearly


304


Π

the influence that a modification in the catalyst preparation may have upon the catalytic reaction itself. Calorimetric Study of Catalytic Reaction. Small doses of the stoi chiometric mixture C O + i 0 are allowed to react with a sample of N i O (250) placed in the calorimeter. A cold trap is placed near the oxide sample. W h e n the reaction of one dose is completed—when no more thermal effect is registered—the final pressure is in all cases low (p < 10" torr). The heat of reaction is plotted in Figure 4A as a function of the total volume of the reaction mixture introduced to the catalyst. Dur ing the reaction of the first doses, the heat of reaction is 78 kcal. per mole. It decreases progressively and, after the reaction of 15 cc. of gas mixture per gram of N i O (250), it remains constant at a value of 69 kcal. per mole. The heat of the homogeneous reaction is 68 kcal. per mole (31). The higher values which are registered ( Figure 4A ) are explained by poison ing of the catalyst by adsorbed carbon dioxide 2


4

CO(g) + i 0 ( g ) = C 0 ( g )

+68 kcal.

C0 (g) = C 0

+29 kcal.

2

2

2

2 ( a d s )

CO(g) + i 0 ( g ) = C 0 2

2 ( a d s )

+97 kcal.

If all carbon dioxide had remained adsorbed on the surface, during the reaction of the first doses the registered heat of reaction should have been 97 kcal. per mole. The initial heats of reaction are lower (78 kcal. per mole). This means that gaseous carbon dioxide is produced during the reaction of the first doses. Therefore, inhibition of the most active sites of the catalyst surface proceeds, but progressively. Further evidence of the self-inhibition of the surface is obtained from the calorimetric results. In Figure 4B the percentage of the heat released during the reaction of doses A , B, and C (Figure 4 A ) is plotted as a function of time. The heat of reaction is produced more slowly for dose C than for Β or A . Therefore, the catalytic activity decreases in the order A > Β > C . The curve corresponding to dose C indicates a steady value of the activity, inhibition of the surface being then completed [heat of reaction, 69 kcal. per mole (Figure 4 A ) compared with 68 kcal. for the enthalpy of the homogeneous process]. These results demonstrate a progressive inhibition of the oxide surface by generated carbon dioxide remaining partially adsorbed. The same conclusion has already been deduced from the study of the surface interactions. For N i O (200), the initial heat of reaction is 90 kcal. per mole (20). Hence, a larger fraction of the carbon dioxide formed remains in the adsorbed state, and the catalytic activity decreases rapidly, after the conversion of increasing amounts of the stoichiometric mixture [Figure 4B: dose A ' after conversion of 0.4 cc. per gram, dose B ' after conversion


48.

EL

SHOBAKY E T

AL.


305

of 5.69 cc. per gram, dose C after conversion of 22.68 cc. per gram. Compare with doses A , Β, and C for N i O (250) in Figure 4 A ] .

®

• Φ

1 8

80-i 68 40·


a no

Figure 4.

15

20

25

cm /g 3

Heats of reaction

A: Calorimetric B: Percentage evolved as a function of time for oxidation of carbon monoxide on NiO(250) for doses A, B, and C (Figure 4A) and Α', B', and C on NiO(200)

The catalytic activity, represented by the rate of evolution of the heat of reaction, is approximately the same for both catalysts when N i O (250) has converted 23.32 cc. per gram of the reaction mixture and NiO(200) has converted 5.69 cc. per gram (Figure 4B, curves C and B ' ) . A n even larger difference of activity is evident from curves C and C , both of which represent a steady value of the activity. The heat of reac tion of doses C and C on both catalysts is close to 68 kcal. per mole (20).


306

OXIDATION O F ORGANIC C O M P O U N D S

II

Carbon dioxide is then desorbed almost entirely to the gas phase, and the coverage of both surfaces by adsorbed C 0 must be close to its maximum value. Since the capacity of adsorption of both catalysts with respect to carbon dioxide is the same (Table II), the difference of their activities cannot be caused by a different coverage of their surfaces by carbon dioxide. A test carried out in a static reactor ( catalyst weight 50 mg., initial pressure of the mixture C O + | 0 3 torr, liquid nitrogen trap to condense C 0 ) confirms that N i O (250) is more active than N i O (200) i n the room-temperature oxidation of carbon monoxide (Figure 5). 2

2


2

Reaction Mechanism. The rate of production of heat as a function of time gives indications of the velocity of the process taking place on the catalyst surface (Figure 4 B ) . For instance, it has been shown (20) that, on N i O ( 2 0 0 ) , the adsorption of oxygen and the formation of CO.^cads) ions are fast processes compared with the adsorption of carbon monoxide or the reaction between C O and COs'uds). From calorimetric results and a kinetic study of the reaction, it has been concluded (8) that the decom position of C0 "(ads) ions by adsorbed carbon monoxide to yield carbon dioxide is the slowest step of the reaction mechanism on N i O (200) (Mechanism I ) . 3

O n the surface of N i O (250), two simultaneous reaction paths are probable. Mechanism I CO(g) = C O

co,.*, + o ( ) 2

C0 8

( a d 8 )

(3)

( a d s )

g

+ N i * = co 3

(ad8)

+ Ni»*

+ CO(g) + N i * = 2 C 0 ( g ) + Ni* +

2

(l) (2)

Mechanism II \ 0 (g)

+ Ni* = 0-

2

0-

( a d s )

( a d s )

+

Ni»

+

+ CO(g) + Ni» = C 0 ( g ) + N i * +

2

(4)

In order to determine which mechanism actually governs the catalytic reaction, it is necessary to compare the rates of the slowest steps of both mechanisms. As in the case of N i O (200), calorimetric experiments show that adsorption of oxygen and formation of C 0 " ) ions (Interaction 1) are fast processes also on the surface of N i O (250). However, the follow ing interactions are slower processes. 3

C0 3

( a d s )

+ CO(g) + Ni»* = 2 C 0 ( g ) + N i *

CO(g) = C O 0-

( a d

s,

( a d S

2

(3)

( a d s )

+ CO(g) + N i

(2)

3 +

= C0 (g) + N i * 2


(4)

48.



307


ο

0

20

Figure 5.

40

60

80

100

rimG(mn)

Kinetic curves for oxidation of carbon monoxide on NiO(200) and NiO(250)

Reactions 2 and 3 belong to Mechanism I, and Reaction 4 belongs to Mechanism II. Therefore, the slowest step of Mechanism II can only be Reaction 4. As has been stated, the calorimetric method does not allow dis crimination between the interactions involving two adsorbed species or one adsorbed species and a gas. Thus, the following interactions would also be proposed as possible slowest steps for Mechanisms I and II. C0 -

( a d s )

0-

) + CO

3

(adS

+ CO

( a d s )

( a d s )

+ Ni** = 2 C 0 ( g ) + N i * 2

+ Ni* = C0 (g) + Ni* +

2

+

(2a) (4a)

Interactions 2 and 4 represent a Rideal mechanism and Interactions 2a and 4a a Langmuir-Hinshelwood mechanism. However, to form C0 "(ads)> Interaction 1 i n Mechanism I, C O must first be adsorbed; since Interaction 1 is a fast process, the adsorption of C O would be the slow step of Mechanism I, and the kinetics of the reaction would depend on pco- However, it has been shown (8, 28) that the reaction is zero order with respect to C O , and therefore the adsorption of C O and its conversion to CO ; r (ads> are faster processes than Interaction 2 which is the rate-limiting step and hence may be written in the form of 2a ( Lang muir-Hinshelwood mechanism ). 3

As for Mechanism II, for which a choice has to be made between rate-determining steps 4 and 4a, the kinetic results (7) show again that the reaction rate does not depend on p. Here again step 4a has to be preferred (Langmuir-Hinshelwood mechanism) to step 4^ (Rideal mechanism ). c


308

OXIDATION O F ORGANIC COMPOUNDS

II

The rates of production of the heat evolved when a dose of carbon monoxide interacts with N i O ( 2 5 0 ) containing either 0" ds) ions (Reac tion 4) or C0 " ads) ions (Reaction 2) are given as a function of time i n Figure 6. In both cases, the same amount of carbon monoxide has been introduced previously to this particular dose. Thermochemical cycles and direct observation of the presence of carbon dioxide i n the cold trap confirm that, during the interaction of this particular dose of C O , carbon dioxide is desorbed to the gas phase. (a


3

(

0

Figure 6.

AO

80

120

160

rime(mn)

Percentage of heat as a function of time for Interactions 2 and 4

The rate of production of heat is greater for Interaction 4 ( Mecha nism II ) than for Interaction 2 ( Mechanism I ). Thus, again the limiting step of both Interactions 2 and 4 cannot be the adsorption of carbon monoxide (Equation 3) since the rates of adsorption of C O , without interaction with preadsorbed species, on nickel oxides containing nearly the same quantity of 0" ds) ions or C0 "(ads) ions should be the same for the same coverage of the surface and the rate of production of heat would be the same. (a

3

Interaction 2a [in the case of N i O ( 2 0 0 ) ] (8) is the slowest step of the reaction mechanism ( I ) . In the case of N i O ( 2 5 0 ) , this mechanism cannot control the reaction rate because Mechanism II is faster (Figure 6), and therefore it prevails. Hence, the most probable mechanism of the room-temperature oxidation of carbon monoxide on N i O (250) is Mecha nism II. Finally, the difference between the catalytic activities of


48.

EL

SHOBAKY E T


AL.

309


N i O ( 2 0 0 ) and N i O (250) is explained by a change of the reaction Mecha nism: Mechanism I on NiO(200) and Mechanism II on N i O ( 2 5 0 ) . Nature of Active Sites. There is no apparent correlation between the increase of catalytic activity and a modification of the electronic structure of nickel oxide, since the electrical properties of both catalysts are identical. It is probable that local modifications of the nickel oxide surface are responsible for the change of its activity and of the reaction mechanism. It should be possible to associate these structural modifica tion with local modifications of the height of the F e r m i level, but it would be difficult to explain the results by the electronic theory of catalysis which considers only collective electrons or holes. A discussion based only on the influence of surface defects seems, therefore, to be more straightforward. The crystal field surrounding the surface ions suggest that the surface of a metallic oxide w i l l be terminated preferentially by anions (JO). However, when nickel oxide is prepared at a low temperature ( 2 0 0 ° C ) , the surface mobility is not large enough to allow a complete reorganiza tion of the surface, and exposed nickel ions may still exist. The increase of the temperature of preparation should result in a more regular struc ture of the outside lattice layer. The existence of a surface mobility i n N i O at 250°C. is demonstrated by the annealing of surface defects in duced at room temperature by neutron bombardment (5). For the nickel oxide used in this work, it has been shown (13) that foreign ions ( L i and G a ) can be incorporated in the outside layers of the lattice, under vacuum, at 2 5 0 ° C , but not at 200°C. These results demonstrate that, at 2 5 0 ° C , the surface mobility may be large enough to facilitate a re organization of the surface layers of the lattice and, consequently, the surface of the nickel oxide prepared at 250°C. should present a smaller number of defects such as exposed nickel ions than the surface of N i O (200). The increase of the temperature of preparation of nickel oxide from 200° to 250°C. therefore produces two types of surface modi fication, both connected with the increased ionic surface mobility—re cession of nickel ions which are more exposed when the oxide is prepared at 200°C. and departure of oxygen with the resulting formation of anionic and cationic vacancies and nickel crystallites. +

3 +

These surface modifications have little influence on the reactivity of the oxide towards the adsorption of C O and C 0 . Both gases must be adsorbed at room temperature on N i ions as has been shown through different experimental results ( 23, 34 ). However, participation of anions in the mechanism of adsorption of both gases is probable since oxygen from either C O or C 0 is exchangeable with lattice anions at room temperature (2,3). 2

2 +

2


310

OXIDATION O F ORGANIC COMPOUNDS

II

In contrast with these results, the increase of the temperature of preparation of the catalyst and the resulting surface modifications influ ence the adsorption of oxygen. Oxygen is adsorbed on nickel ions, and the contributions of the anions to this adsorption may be only negative. The low surface coverage b y adsorbed oxygen suggests that i n both cases the adsorption sites are lattice defects with exposed nickel ions (5). Moreover, for a dissociative adsorption to occur, pairs of sites must exist. However, polarization of the oxygen molecule adsorbed on an active site may induce a second adsorption center for oxygen ( 5 ) . This seems to be the case for the adsorption of oxygen on anionic vacancies which exist i n N i O (250) since sequence 0 — C O has shown that the interaction between the most energetically adsorbed oxygen ions—i.e., probably adsorbed on anionic vacancies—and carbon monoxide yields CO.^ads), this species being formed by reaction of one carbon monoxide molecule with two oxygen ions adsorbed therefore on adjacent sites. This interaction is not registered on NiO(200) where only C0 (ads) is formed.


2

2

O n the less active sites of N i O (250) and on N i O (200), the inter action between oxygen adsorbed probably on exposed nickel ions and carbon monoxide. yields carbon dioxide directly. Hence, the mean dis tance between such sites must be short enough to allow the adsorption of oxygen with dissociation of the molecule but long enough to restrict the interaction of C O to one oxygen ion. When the temperature of prepa ration of N i O is increased from 200° to 2 5 0 ° C , the number of these sites and their average activity decrease as a consequence of the recession of nickel ions. Desorption of carbon dioxide, formed by the interaction between oxygen ions adsorbed on this type of sites and carbon monoxide, is then possible on N i O (250) but not on N i O (200) and therefore gives a new path to the reaction mechanism ( II ). This desorption of C 0 must be a cooperative phenomenon with activation of the surface intermediate complex by the energy of interaction between 0~ ds) and C O because direct adsorption of carbon dioxide on these sites is irreversible at room temperature (28, 30). 2

(a

Conclusions The increase of the catalytic activity i n the room-temperature oxida tion of carbon monoxide, which results from the increase of the tempera ture of preparation of N i O from 200° to 2 5 0 ° C , is related to the difference in the reactivity of oxygen adsorbed on both surfaces. The interaction between adsorbed oxygen and carbon monoxide has roughly the same velocity on both oxides. But on N i O (200) this interaction yields only 2 (ads)> whereas on NiO(250) the same interaction produces C O 3 (ads) on the most active sites (anionic vacancies) and C 0 ( g ) on the less

co

2


48.



311

reactive sites ( partially recessed nickel ions ). Hence, the same interaction causes the inhibition of the surface sites on N i O (200) but yields the product of the catalytic reaction C 0 ( g ) for N i O ( 2 5 0 ) . Since this inter action is the limiting step of the reaction mechanism (Mechanism I I ) on N i O (250), whereas on N i O (200) the limiting step is the decomposi tion of C0 " ads) b y CO ads) ('Mechanism I ) , it is concluded that the catalytic activity may be governed by local chemical defects and not only by collective properties. Therefore, the activity may be correlated with the chemical structure of heterogeneous surfaces, and more specifi cally i n this work, with the energy spectrum for the adsorption of oxygen. 2


3

(

(

Literature Cited (1) (2) (3) (4) (5) (6) (7) (8) (9) (10) (11) (12) (13) (14) (15)

Arghiropoulos, B., Teichner, S. J., J. Catalysis 3, 477 (1964). Bailly, J.C.,Gravelle, P.C.,Teichner, S. J., Bull. Soc. Chim. 1967, 1620. Bailly,J.C.,Teichner, S. J., Bull. Soc. Chim. 1967, 2376. Bielanski, Α., Deren, T., Haber, T., Sloczynski, T., Trans. Faraday Soc. 58, 166 (1962). Charman, H. B., Dell, R. M., Teale, S. S., Trans. Faraday Soc. 59, 469 (1963). Cimino, Α., Molinari, E., Romeo, G., Z. Physik. Chem. 16 (NF), 101 (1958). Coue, J., Ph.D. Thesis 149, Lyon, 1963. Coue, J., Gravelle, P. C., Ranc, R. E., Rue, P., Teichner, S.J.,"Proceed ings of 3rd International Congress on Catalysis," p. 748, North Holland Publishing Co., 1965. Dell, R. M., Stone, F. S., Trans. Faraday Soc. 50, 501 (1954). Dowden, D. Α., Wells, D., "2nd International Congress on Catalysis," p. 1499, Technip, Paris, 1961. Dry, M. E., Stone, F. S., Discussions Faraday Soc. 28, 192 (1959). El Shobaky, G., Ph.D. Thesis 403, Lyon, 1966. El Shobaky, G., Gravelle, P. C., Teichner, S. J., Bull. Soc. Chim. 1967, 3244, 3251, 3670. El Shobaky, G., Gravelle, P. C., Teichner, S. J., "International Congress on Calorimetry," p. 175, CNRS, Marseille, 1965. El Shobaky, G., Gravelle, P. C., Teichner, S. J., Trambouze, Y., Turlier, P., J. Chim. Phys. 64, 310 (1967).

(16) Garner, W. E., Stone, F. S., Tiley, P. F., Proc. Roy. Soc. (London) A197, 294 (1949). (17) Ibid., A211, 472 (1952). (18) Gravelle, P. C., J. Chim. Phys. 61, 455 (1964). (19) Gravelle, P. C., El Shobaky, G., Urbain, H., Compt. Rend. 262, 549 (1966). (20) Gravelle, P. C., Teichner, S. J., J. Chim. Phys. 61, 527, 533, 625, 1089 (1964). (21)

Keier, N. P., Roginsky, S. Z., Sazonova, I. S., Dokl. Akad. Nauk USSR

(22)

Keier, N. P., Roginsky, S. Z., Sazonova, I. S., Izvest. Akad. Nauk USSR

106, 859 (1956). Ser. Fiz. 21, 183 (1957). (23) Marcellini, R. P., Ranc, R. E., Teichner, S. J., "2nd International Congress on Catalysis," p. 289, Technip., Paris, 1961. (24) Marcellini, R. P., Teichner, S. J., J. Chim. Phys. 58, 625 (1961). Mayo; Oxidation of Organic Compounds Advances in Chemistry; American Chemical Society: Washington, DC, 1968.

312 (25) (26) (27) (28) (29)


(30) (31) (32) (33) (34) (35) (36) (37)

OXIDATION O F O R G A N I C

COMPOUNDS—II

Merlin, Α., Teichner, S. J., Compt. Rend. 236, 1892 (1953). Parravano, G., J. Am. Chem. Soc. 75, 1448, 1452 (1953). Parravano, G., Boudart, M., Advan. Catalysis 7, 47 (1955). Ranc, R. E., Teichner, S. J., Bull. Soc. Chim. 1967, 1717, 1730. Roginsky, S. Z., Tselinskaya, T. F., Zh. Fiz. Khim. USSR 22, 1360 (1948). Rue, P., Teichner, S. J., Bull. Soc. Chim. 1964, 2791. "Selected Values of Chemical Thermodynamic Properties," Nat. Bur. Std. Bull. 500 (1952). Schwab, G. M., Block, J. Ζ. Physik. Chem. 1 (NF), 42 (1954). Schwab, G. M., Block, J., J. Chim. Phys. 51, 664 (1954). Teichner, S. J., Marcellini, R. P., Rue, P., Advan. Catalysis 9, 458 (1957). Teichner, S. J., Morrison, J. Α., Trans. Faraday Soc. 51, 961 (1955). Wagner, C., Hauffe, Κ., Z. Elektrochem. 44, 72 (1938). Winter, E. R. S., J. Catalysis 6, 35 (1966).

RECEIVED October 9, 1967.


Mechanisms of the Room Temperature Oxidation of Carbon Monoxide

Recommend Documents