Chapter III: The Structure and Physico-Chemical Properties of Heterogeneous Catalysts for Reactions Involving Molecular Oxygen

41

Chapter

III

THE STRUCTURE AND PHYSICOCHEMICAL PROPERTIES OF HETEROGENEOUS CATALYSTS FOR REACTIONS INVOLVING MOLECULAR OXYGEN Since the rates of catalytic processes are essentially determined by the chemical composition and structure of the catalysts, it is advisable to consider first of all the catalysts for the reactions to be considered. Processes involving O2 may be accelerated by transition metals, their alloys, metal oxides, carbides and by some other substances. Metal and Metal-Like Catalysts Metals Structure and Physicochemical Properties A peculiar feature of metals in the condensed state is that they have free electrons which give rise to high electrical and thermal conductivities, a metallic lustre, a positive temperature coefficient of electrical resistance, plasticity, etc. In the formation of a solid crystal from isolated metal atoms, the valence electrons cease to be bonded to separate atoms and form an electronic gas into which the skeleton of positive metal ions is immersed. The gas of "collectivized" electrons compensates for repulsion between the ions. Thus, a special type of chemical bond, the metallic bond, appears. When the usual covalent bond is formed between two atoms to give a molecule, unpaired electrons are also collectivized but electron pairs remain localized at the atoms; in the case of metals, the collectivized electrons are not localized and belong to the whole crystal. Metal crystals are formed from atoms of both transition and nontransition elements. The latter belong to the main subgroups of I-IV groups of the Mendeleev Periodic Table. The structure and physicochemical properties of metals have been well described / 1-4 /; Metal crystals usually have highly symmetric structures with high coordination numbers. In most cases, the structure corresponds to a compact packing (a face-centered cubic lattice or a

42

hexagonal lattice with a coordination number equal to 12). In other cases, either a simple body-centered cubic structure or a tetragonal one are formed. Only some metals (for example, manganese) give more complex lattices. Under various conditions (i.e. temperature, pressure), the metallic element can form several polymorphic modifications with different crystal structures. Some of these may not display metal properties (e.g. grey tin) while some polymorphic modifications formed by nonmetallic elements (e.g. graphite) have metallic properties. A significant characteristic of metals is the electronic work function, ~,which usually varies from 2 to 5.5 eVe These values are considerably lower than the ionization potentials of atoms and this suggests that there is an increase in the kinetic energy of the electrons on going from isolated atoms to metal crystals. For a description of the electronic structure of metals, the band theory is used. It assumes that the ionic skeleton of the crystal within the metal causes an effective periodic field in which collectivized electrons move. Each of the electrons is characterized by a quasi-pulse, p (similar to pulse of a particle in free space) and a discrete parameter n. The corresponding wavefunction is similar to that of the free electron. This leads to a periodic dependence of the energy on n, and to a band pattern of the energetic spectrum of the electrons. In the formation of a crystal,the levels of the valence electrons of N isolated metal atoms form corresponding bands, each of them containing N levels (2N states). In real cases, N is very high, so that the distances between the levels in the band are small and, within the band, the energy is changed almost continuously. The deeper the electron level in an atom, the narrower is the corresponding band. The latter are practically all formed from the external valence electrons, while the electrons of deep levels remain localized at the atomic nuclei. In a three-dimensional space of pulses, the bands are represented by polyhedra covering each other(the Brilluene bands). They can overlap but they keep their individuality. The electrons in the metal occupy the energy levels in such a way as to reach a minimal total energy as well as obeying the Pauli principle.Levels of valence electrons are incompletely filled and this is peculiar to metals. By elevation of the temperature or by optical excitation, the electrons may occupy vacant levels. The

43

°

energy of the top occupied level at K (the Fermi level) corresponds to work function, ~. At T /0 K, the Fermi level is that which can be occupied with a probability of ~ • The band theory only leads to adequate results in the calculation of the energies of crystal lattices, interatomic distances, etc., for the case of nontransition metals, particularly the alkali metals. For the transition metals, the Pauling theory /5/ is sometimes used. According to this, among the valence electrons, one should distinguish between binding and atomic electrons. The former are described by hybrid spd-functions like those in coordination compounds. They correspond to occupied levels in the complex spd-bands of the transition metal. The binding electrons form covalent bonds between the atoms and determine the strength and structure of the crystal lattice. Atomic electrons described by d-functions do not contribute to metal-metal bonds but determine the paramagnetism of the metal. Calculations of the weights of d-states in metal bonds are based on the assumptions of the combination of different electronic structures of the metal atom to give agreement between the estimated and observed magnetic properties. The use of metals as catalysts depends upon their thermal stability, i.e. the strength of their crystal lattices. In this respect, transition metals are essentially superior to nontransition metals since the sublimation heats and melting points for the former are much higher than those for the latter. However, there are among the nontransition metals some quite thermally stable solids (e.g. AI). The fact that these are not used in catalysis suggests that other factors are of great importance. Under the conditions of catalytic reactions involving 02' the metals can be influenced by the reaction mixture. In such reactions as homomolecular oxygen exchange, the main effect of this sort is the oxidation of the metal into an oxide phase (or oxygen dissolution in the metal) while, during the catalytic oxidation processes, the formation of polymer films on the sUrface, etc. is also possible. In other words, chemical stability of the metal is also essential. The most common and most reactive participant of all the reactions discussed is molecular oxygen. Hence, the resistance of the metals to oxidation to the bulk oxides is of primary importance. Table 2 presents the data on standard enthalpy changes in the reactions

44

TABLE 2 The Heats of Formation of the Lower Metal Oxides (- ~ ~98' kcal(g-at 0)-1) and Initial Heats of 02 Chemisorption on Metals (q~, kcal(g-at 0)-1) Metal

Oxide

-.1 H0298

0

qs

Metal

Oxide

Ir Pd Hg

I Rh

Ru Os Cu

Tl Bi Re Pb Sb Co Ni Cd Fe Ge Sn

w Mo In Cs

AU203 Ag20 PtO

°7 17

Ir20 3 PdO Hg 20 Rh20 Ru02 Os02 Cu20

20 21 22 23 28,36 31 40

CuO T120 BiO Re02 PbO Sb20 3 CoO NiO CdO FeO GeO SnO W0 2 Mo0 2 In20 3 Cs20

37 42 50 51 52 56 57 58 61 64 67±6 68 68 66;71 74 76

* Integral value

H~98

0

qs

/7,8/

/7,8/

Au Ag Pt

- L1

55 /22/ 63 /22/ 15-35 /21,23/

-

34*/21/

-

38 * /21/

-

57 /28/; 60 /22/

-

-

-

-

50 /21/ 54 /21/

-

68 /21/ 66 /25/

-

97 /21/ 86 /21/

-

Ga Zn K Mn Nb Ta Na

Ga20 ZnO K20 MnO NbO Ta20 5 Na20

V

vo

Cr Ti Gd

CrO TiO Gd 20

u

uo

Zr Al Ba Sc Sr Li Be La

zr0 2

Mg

Sm Tb Th Pr Ca y

Nd Ce

3

A1 20 3 BaO Sc 20 3 srO Li 20 BeO La 20 3 MgO Sm20 3 Tb20 3 Th02 Pr20 3 CaO Y203 NdO CeO

82 83 87 92 97 98 99 100 105 124 126 129 131 133 133 137 141 142 143 143 144 145 145 145 146 152 152 167 178

-

-

75 /211 104 /211 106 /211

87 /21/ 118 /21/

-

-

106 /21/

-

-

-

-

-

45

m Me (solid) + n

1

2

02

(III.1) •

Using this parameter, the metals are conventionally distributed between three groups: 1. Metals which form stable oxides ( -,1 H~98> 65 kcal(g-atom 0)':-1 These are the alkali and alkali earth metals;metals of the subgroups containing Sc, Ti, V, Cr and also Mn; and the rare earth metals, the actinides, Zn, AI, Ga, In, Ge and Sn. 2. Metals which form oxides of moderate stability ( -L1~98 = 40 - 65 kcal(g-atom 0)-1). This group covers the transition metals, Fe, Co, Ni, Cu, Re, and the nontransition metals, Cd, Sb, Pb, Bi, Tl. J. Metals which form relatively unstable oxides ( - ,1 HO 298 <: =

< 40

kcal (g-atom 0)-1). These are the platinum metals, Ru, Rh, Pd, Rh, Ir and pt, and also Ag, Au and Hg. From this point of view,one should expect that the third group of metals will remain unchanged during oxidation. Under definite conditions ( i.e. great excesses of reducing agents), the second group of metals can also be used. In practice, only transition metals, mainly noble ones belonging to the groups mentioned, are used in oxidative catalysis. Some of the nontransition metals (for example, mercury) do not satisfy the requirement of thermal stability, but the other ones might be used as catalysts. The fact that they are not used suggests that the thermal and chemical stability are only restrictive factors; their ability to give definite intermediate chemical interaction with the reagents is of primary importance. Transition metals of the third group above, as follows from practice, satisfy all the requirements formulated. Metal catalysts are used in the form of compact samples (wire, gauze, foil, plate, powder, film or single crystal) or are supported on various carriers (corrundum, alumina, silica, pumice, coal, etc.). During catalytic reactions involving 02 metal catalysts can undergo various changes. Highly dispersed metals can be sintered if the temperature of their preparation had been lower than that of the catalytic reaction. For such catalysts as Ni, Fe, Co or Cu, great excesses of the oxidized substances are required to avoid the bulk oxidation of the metals /9/. At high temperatures,the oxidation of such me-

46

tals as Pt, Pd, Os and Ir takes place, for example, in the oxidation of ammonia /10/ or NH + CH mixtures /11/ at 1000 0C. The 3 4 combined influence of temperature and the reaction mixture leads to catalytic corrosion of Pt - Rh gauzes. Noble metals can dissolve many monolayers of oxygen /12, 13/. In the oxidation of organic sUbstances, metal surfaces are covered with stable polymeric films or organic residues /14/. The reaction mixture generally levels out any initial differences in metal catalysts which have the same chemical composition but had been prepared in different ways. This causes approximate constancy of their specific catalytic activities /15/. The Thermodynamic Characteristics of Surface Compounds Formed by Metals The rates of catalytic reactions depend upon the energies of the bonds between the reagents and the catalysts. In the case of simple gases (e.g. 02' H2 , N2 ) , these energies coincide with the heats of adsorption, q. For the measurement of differential adsorption heats, qa' at various surface coverages, 8, calorimetric methods /16, 17/ are used. For instance, the heats of formation of surface oxides:

12

°2

(111.2)

+ (

are measured in this way. The values of qe can either be constant or decrease with increasing B. Calorimetry is sometimes used to obtain heats of adsorption indirectly from another surface reaction at equilibrium /18/. For example, by measuring the values of the heat, q', of surface reduction: (111.3)

one can find qs = q" H2 +

1

2

- q', where q"

is the heat of the reaction

02 = H2 0 . Instead of (111.3), the reaction

CO + (0) = CO 2 + ( can be used.

)

47

Since the equilibrium constant of the reaction (111.2) at fixed

8 is

Ka

=

- 1

Po 2, one can calculate qs values by measuring the 2

temperature dependence of Ka (at B = const) and using the Clausius - Clapeyron equation /16/: d In KEf qs d (

*)

--R

what leads to Fig. 3 - Initial heats of oxygen chemisorption on metals against heats of formation of lower metal oxides (see Table 2)

Ti Ta

Wo Nb 8 oMO oCr

/00

CI/

fie

Mn 0

fie

aNi Co

o

o '----~-----l..----l.

50

/00

/50

I

-11H!Da /xCflt (q-fltOF

1 In Po = 2

2

qs - -- + RT

conat ,

(III.6)

When the adsorption isotherms of the gas at various values of T are known, one can draw a line at given 8 to find the equilibrium Po values at constant 8 • 2

The equilibrium constants K of the surface equilibria such as (111.3) and (111.4) can be measured directly /19, 20/. The values of q are then calculated using the equation: (III.1) where the standard entropy (Ll So) and the Gibbs free energy (L1 GO) changes refer to (III.3) or (III.4); the LI SO value is estimated independently, while Ll GO = - RTlnK. For reactions involving O2, the values of qs are of primary importance. One can summarize the following trends. 1. The initial heats of oxygen chemisorption, q~, are approximately proportional to the heats of formation of bulk oxides, - L1 H~98 /21,24-21/ (see Fig. 3). Usually, q~ > -LlH~98 but

48

o sometimes qs

0

- L1 H298• to reaction (111.1) increase 2. Values of - L1H~98' referring the metal (Fig. 4). with decreasing oxidation state of ~

Fig. 4 - The dependence of the heats of reaction (111.1) on n/m (thermochemical data /7,8/): 1 - Ce, 2 - Cr, 3 - Mn, 4 - Ge, 5 - Fe, 6 - Co, 7 - Bi, 8 - Cu, 9 - Hg, 10- Ag, 11 - Mo.

fjO/!
Ttl Wo

100

Nb 8

cr

Mo 0 ........"'2 o ., Fe

•v: ., 8",1 .,

50

.../fTa

Fig. 5 - Initial heats of chemisorption of various gases on metal films /25/ against the heats of formation of the lower metal oxides-

Co ....PW ., 6fe Ta Cr Nig W ~ 00 Mo

H 2

Fe

50

100 150 -I1Hf;o /ircaL(q-atO;-1

It is advisable to compare the values of q~

with tho se of - L1 ~98

for the lower oxides since at small values of 8, the lower surface oxides are formed (see,for example, /18/). Similarly has been shown /24, 25, 27/ that the heats of chemisorption of H2, N2, etc. are correlated with the heats of formation of the hydrides, nitrides, etc. 3. The heats of chemisorption of various gases are approximately correlated with the heats of sublimation of the metals /24/.

49

4. The heats of chemisorption of various gases ly proportional to the heats of formation of the g-atom of oxygen /30/ (see Fig. 5). 5. The two last trends suggest that the values ous gases are approximately proportional to each ven metal, the qO values usually decrease in the

are approximatebulk oxides per of qO for variother. On a giorder /25/: (III.8)

6. The bond energy of the given atom A with the surface has a value of the same order as the bond energy of A with the isolated metal atom. Fig. 6 presents the correlation between the heats of the processes: Me(gas) + A(gas) = MeA(gas) + e Me(solid) + A (gas) = Me(solid) A (ads) + s' (A 0, N, H), where

=

£'=

q

O

1

+ -2 DA

2

(DA is the dissociation energy of A2). 2 e/KcaLfg-atr l 200

Fig. 6 - The dependence of the bond energies of atoms of oxygen (0), hydrogen ( ~ ), nitrogen ( 0 ) with metal surfaces, s' , upon the energies of the corresponding bonds in the gaseous molecules, c (values of t are taken from /31/).

150 100

50 50

100

150

200

250

e'/KcaL (g-tItt'

One can

imagine the cycle:

(Me) s = Me(gas) + (

)s - L

(III.10)

D

(III.11)

i A2(gas) = A(gas) - 12

A2 Me(gas) + A(gas) = MeA(gas) +

e

(III.12)

50

= (MeA)s + L' A2(gas) = (MeA)s + qO

MeA(gas) + ( (Me)s +

t

)s

where (Me)s is metal atom in the surface layer, and )s is a vacancy corresponding to this atom. On the basis of the above cycle: qO

=e -

1

2

DA + L' - L 2

(III.14)

or taking Eq.(III.9) into account: Sf =

C

+

L

-

L'

It the Me - A bond does not affect the bond energy of the Me atom with the neighbouring metal atoms, L

~

(III.16)

L'

and we arrive at:

e ;::: e

(III. 17)

which can explain the rough regularity shown in Fig. 6. It tollows trom these arguments that the main contribution to the energy of chemisorption is given by the energy e of the local chemical interaction between A and Me. A similar cycle may be written for the formation of the bulk compound MeA(solid): Me(solid) +

t A2(gas) = MeA(solid)

(III.18)

which leads to:

where

~A

and

~e

are sublimation heats. Combining Eqs (111.19)

and (111.14) one arrives at: (III.20) or (III.21)

51

(II1.22) A .... :.:~ear wr:.Gn A

r-eLa t Lon sh.l p be twe en qO and

will be observed

.Ls cone t ant or is linearly correlated with qO. Theoreti-

cu I, a.rl:J.]ysis of A

If

-AH~98

Olle 3c~epts

is difficult.

Eq.(III.16),

I. is similur toLlIIe while l' is similar to LrlieA' a com:)FlIGation is possible which leads (on the basis of Eq, (111.22» JL:l~e

to

~A"""'O'd -..- . an

q 0""", ,.....

ATJo -~'.298.

T'w ill - 0 bo nd er.·orgies in molecules of MeO are usually higher than the :\le - ;{ bond ene rgLe s in MeH /31/ which explains the high-

er

cheraisorp~ion

(P::Lg.

5).

heats observed. for 02 in comparison with H2

The initial hea.ts of chemisorption of gases on metal~ have been cstiillated using equations similar to the Pauling equation /32/ for the bond energy of AB:

',lthere Xi are electronegativities. To calculate the bond energy of A with the sUl~ace, similar relations /16, 33, 34/ were used; for instance:

or ta,;,;irlG into account Eq.(III.9):

(III.26)

(if

hl work f unc t Lon }, one can obtain satisfactory agreement

between

calculated and. expez-Lmen taL values of the heats of chend scz-p t Lon on transition metals /34/. With 02 and N , 2 liscre;Jd,ncies occur , since the model is imperfect.

'1.y.~rO&e"L

1;:110

52

Equations such as (111.26) provide some arguments in favour of correlations between heats of chemisorption and heats of sublimation of metals. Since Me - 0 bond energies in molecules of MeO are of primary importance in elucidating trends concerning the heats of oxygen chemisorption t quantum-chemical calculations of the Me - 0 bond energies (Me = Tit Vt Cr, Mn t Fe t Co and Ni) were made /35/ using the expanded Huckel method /36/. Resonance integrals were calculated using the Cusachs formula /37/ and for the basic atomic orbitals t the Slater functions /38/ were accepted; partial and total electron populations on atoms and bonds were calculated using the Mulliken equations /39/. Orbital ionization potentials were taken from references /43/ and /44/ which lead to similar results. The program has been described in references /40/ and /41/. The calculated potential energy curve (near the minimum) for Fe - 0 is presented in Fig. 7. Table 3 shows equilibrium distances found theoretically for all the systems considered. In the case of TiO, VO and erO t the values agree with experimental data /42/. Fig. 8 presents the calculated and experimental energies of

o

1.0 1./ 1.2 U 1.4 !.5 1.8 1.7 1.8 I.g ny,40

-/

-2 -J

-4 -5

-6

[lev Fig. 7 - Calculated potential energy curve for Fe o /35/. Fig. 8 - Comparison of calculated Me - 0 bond energies (1) with experimental values (2) and energies of oxygen chemisorption on metals (3).

t ctllcleV

!4

12

lex/J.

/0 8 6

Iwrl(q-at;-I 160 150 140 lJO 120

4 2

o

//0

L...-_--.,,;:b-_ _....JIOO

f,'exfl:llfcat(;-rrtFI

Ili 1611 /50 /40 /30

/20 //0 Ti V CrMnfeCoNi

53

Me - 0 bonds as well as the experimental energies of oxygen che-

misorption on metals,

f

s Me-O

0

= qs +

2'1

0

D ' where qs are initial

02

heats of adsorption /21/ (for Co the values q~ = 61 kcal(g-atom 0)-1 /45/ is also taken). All the values compared change in parallel to one another. Similar results were obtained by the valence bond method /46/. Using the expanded Huckel method, the same correlation was found to be valid for the bond energies of metal with H, Nand CO /47/. TABLE 3 Calculated Equilibrium Distances and Dissociation Energies for M - 0 /35/ Metal

Equilibrium distance/ K

Dissociation

Metal

energy /ev calc.

Calcu- Expelation riment

Calcu- Expelation riment Ti V Cr lin

1.5 1.5 1.5 1.6

1.62 1.89 1.63

13.602 12.994 12.706 9.576

Equilibrium distance/ A

Fe Co Hi

1.6 1.8 1.9

Dissociation energy leV (calc.)

6.405 2.855 1.164

There is definite trend to q~ decreasing with an increase in work rune tion 'f (Fig. 9). It can be shown that

Fig. 9 - Comparison of initial heats of oxygen chemisorption on metals with the values /48/ of work function

OL..-----.l..

4

-l-_

54

W - 2

1

Cf

- (2" DO

- E

/I

(III.28)

)

2

where En is the energy of the addition of two electrons to 0 to form 0 2-; W is the energy of 0 2- addition to the metal which had lost two electrons (~DO - E 1/ = coria t }; If changes in W do not 2

exceed those in ~, one should expect a decrease in q~ with increasing cP • Similar conclusion can be made on the basis of the relation:

r -

O

qs = W' -

1 DO 2

(2

- E

I

)

where the cycle involves 0- ions (one-electron transfer). The term

E W' = qO

s

+

I

~

= 25

kcal(g-atom)/31/. Then

Cf

+ 25.

Using the last equation, one finds that WI is 174 kcal(g-atom)-1 for Pd and 211 kcal(g-atom)-1 for Mo, i.e. changes in W' are of the same order as those in ~. This suggests that the correlation q~ - ~ is only approximate. Dependence of Adsorption Heats Upon Surface Coverage Adsorption heats are known either to be constant or to decrease with increasing 8 116, 25/. The former corresponds to an ideal adsorbed layer, the latter to a real one 1491. The heats of chemisorption of oxygen, qs (Fig. 10) on the films of Co, Ni, Nb, Ta, Fe and Ti are constant up to a monolayer while, on Cr, Mo, Wand Mn, the value of qs starts to decrease at & = = 0.5 - 0.8. Almost all the metals mentioned can adsorb several monolayers of oxygen. On Pt and Pd, the values of qs decrease with increasing 8121/. The initial heat of oxygen chemisorption on Pt black at 70C is equal to 33 kcal(g-atom 0)-1 which agrees with other data (Table 2). The values of qs then decreases slightly (qs = 28 - 31 at &= = 0.05 - 0.85, qs = 26 kcal(g-atom 0)-1 at 8 = 0.9 150/). The changes of qs with 0 are presented for Cu, Ag and Au in Fig. 11 (see also I 28, 51, 521). In the case of Cu, bulk oxida-

55

tion takes place at high B. According to reference /51/, the surface oxide (Ag20)s is formed at small 9 on Ag, the qs values being at a maximum. The / 0 ............ oxide structure is Ag Ag, the degree of oxidation of the silver being + 1, and the oxygen being charged negatively. At higher 9 , other oxides are formed. There may be two species which correspond to a Ag : ratio of one. The first involves atomic anions of oxygen, i.e. (Ag202)s' the oxide of Ag+ and Ag3+; the se-

°

flS/KCfIL(g.atrl

125~"".,

2

/00

/ J

75 55

~~

100 ~~:;:;::~ 75~."...."....-"",,-

s

o

50~~~-4.,~

00

o 1 2 o a» /

1

"-

1/

5 25'---~---l-.:l""""!....-L----:-~

I

,}

1.5

4(1.4) 2(2.3,5)

o o

0.25

il/

s» a:

0.75

1 (6,7,9) O,J il4 (8/0/1)

a

Fig. 10 - The dependence of the heats of oxygen chemisorption, qs' on metal films on the amounts of adsorbed oxygen, a (ml NTP per 100 mg of the film)/2Y: 1 - Ti, 2 - Ta, 3 - Mo, 4 - Cr, 5 - Fe, 6 - Nb, 7 - W, 8 - Mn, 9 Ni , 1 Co, 11 -

Rh

°-

fls/KCtrt(;-ut;-1

~~,I

Fig. 11 - The dependence of qs ona for CU, Ag and Au ( a is the-amount o~ chemisorbed-ox~gen, ml NTP per m ): 1 - Au (120 0 ) , 2 - Ag (110 0 ) , 3 - Cu (110-115 0 /22/

~t~2 :~" "l

- I--,---I,:: /

tl,/O

)

I

0.20 '0.25/,05 a

cond contains molecular oxygen of the peroxide type. Since the heat of the conversion of (Ag20)s into (Ag202)s is lower than that of the (Ag) s transformation of (Ag) s into (Ag20) s , the value of qs decreases when 9 increases. At high 8, the highest oxide, (Ag203)s is formed with the minimum energy of chemisorption. A

56

similar pattern is thought to be valid for gold /22/. The above mentioned calorimetric qs values are in general accordance with the data obtained by the method of surface reaction equilibrium (see Eq. (111.3»/20, 54-56/, the exception being Pt. Several reasons have been given for a decrease in adsorption heats with increasing 11 /16, 25/. One of them is energetic nonuniformity of the surface. They can be caused by the existence of different crystal faces, edges and corners each with different adsorption potentials. When the distribution of surface sites with adsorption heats is continuous, one can distinguish two main cases /65-67/; 1) an evenly-nonuniform surface for which adsorption heats decrease linearly with increasing 8; 2) an exponentially-nonuniform surface for which adsorption heats decrease linearly with increasing In 11. Another reason is mutual repulsion of equally charged adsorbed particles. Several physical models describing the last effect have been proposed. According to references /68/and/69/, the adsorbed layer is considered as a charged condenser. This model leads to an overestimation of the electrostatic effect /70/ and is insufficient to explain the experimental dependence of qs on 8. An analogy /71/ between the effect of the mutual influence of adsorbed species and the effect of substituents in aromatic compounds has been discussed. The corresponding model of surface electronic gas has been developed /57, 72/. The theory explains a linear decrease in q with growing 11 for metals. Another reason for changes in qs with B is associated with different forms of the adsorbed species formed from the same substance. For example, in the case of oxygen chemisorption, several surface oxides, (MemOn)s' can exist with different stoichiometries and stabilities. The oxides can differ in the nature of the chemical bonds formed. Hence, (MeO)s may involve atomic anions of oxygen, while (Me02)s may include peroxide (-0-0-) groups. One can assume that on increasing the value of n/m in a range of (MemOn)s oxides, the form of oxygen remains constant but the degree of metal oxidation increases. Such an approach was used /22, 51/ for 02/Ag (see above). By analogy with the bulk oxides (Fig.4), one can believe that when n/m increases, the heat of formation of the oxide either remains constant or decreases, which explains the observed relationship between qs and 11.

57

Adsorption heats can depend on 8 even when the form of the adsorbed species is kept constant if there is a mutual influence of adsorbed particles one with another. Relation (111.15) takes this possibility into account since L and L' depend on 8 (in the general case). Approximations (III.16)and(III.17) refer to small values of e when the above effect is negligible.

J ... l~ .-atom Fe o-atomO

• • • • • • • • •

Fig. 12 - A chain of metal atoms with adsorbed oxygen species

In order to discriminate between the effect of the mutual influence of adsorbed species and other effects, quantum-mechanical calculations are likely to be useful. In most cases a one-dimensional chain of atoms of the crystal was used as a model of the surface /58, 59/. However, even for such a simple model, analytic results can be obtained only when each atom is described by one orbital. (The calculation using these assumptions was made /60/ for a three-dimensional homoatomic crystal with a simple cubic lattice). The above simplification seems unlikely to be correct. We have attempted to calculate the electronic structure of a long chain of atoms with a real set of orbitals, the chain being connected to two chemisorbed atoms (Fig. 12). The calculation was made exactly as in the case of a large molecule. This approach allows one to interpret the results using the band theory. A theoretical analysis of oxygen chemisorption on iron was made. We calculated one-electron functions and energy levels of the chain containing 15 atoms of Fe, and then examined the changes in one-electron spectrum caused by adsorption of one and two oxygen atoms. The expanded Huckel method was used /41/. The Fe-Fe distance was equal to that in a crystal /61/, the Fe-O distance being equal to that in the FeO molecule /35/. Initially, the electronic structure of two subsystems (metal chains consisting of 11 atoms and a molecule of FeO) was considered. The energy levels of the chain form four groups of independent states. Three of them give simple bands corresponding to Jdxy-' 3dx z- and 3dyz- orbitals, the fourth one consisting of 4s-, 3dz2-

58

["f/!J.5eV

Fig. 13 - Energy levels of the 3dxy-' 3dxz- and 3dyz-

Fig. 14 - Energy levels of the 3dz 2 -,3dx 2 -y 2 - and 4s -

bands of the chain

bands of the chain

and 3d 2 2 - orbitals. The energy levels in the first two bands x -y are descr~bed by the formula: Ie = 1Cn

(n= f, Z•...

#+1

t¥ )

where N is a number of atoms in the chain and ~ is the resonance integral of interaction of neighbouring atoms. Errors in calculating the bands boundaries do not exceed 4-5% (0.004-0.05 eV). The electronic levels of the bands are presented in Fig. 13. The wave functions of their levels are expressed by

~k

f =~

[sin

kn

~i

(r)

(III.31 )

where n is the number of the atom in the chain, ~ ni is the corresponding atomic orbital. The functions (111.31) are delocalized over the whole chain. The three sub-bands formed by the 3dz2-' 3dX2_y2 - and 4s - orbitals are divided into the 4s-band and two overlapped 3d-bands, the 4s - band and 3d - bands being separated by an energy gap (Fig. 14). The one-electron energy levels of the overlapping .bands are expressed by the equations

59

(III032) is the ionization potential of the 3d-orbitals; and~aa Jd and flab are resonance integrals of the 3d-orbital interaction for neighbouring atoms (3 d z2 with 3dz2 and 3dX2_y2).

where I

The energy diagram for FeO is presented in Fig. 15 (the axis of the molecule coincides with the Z-axis). In the ground state, two lower r3 -states, the lower :J!x-' jfy - states and the nonbinding

Fe-O 0 4 (.10)

o

0/11.75} Jlzx,Jlzv (/4..5)

C'\t=l5.5

t=/8.52 2sFig. 15 - Energy diagram for FeO

~

E=/!J.85(J!J.5)eY

Fig. 16 - Scheme of 3dxy2Px-(3dyz- 2Py) - states of the chain with one adsorbed atom

3d _ y2 x2

and 3dxy - states are completely filled; the upper :J!x-

- levels are half filled. The :Jlx ' 11y - stat es are linear combinations of the 3dxz -, the 2px - and 3dyz- , and the 2py -

and

jf y

orbitals respectively. The 6 -states are linear combinations of the 3d z 2 -, 4s -, 2p z - and 2s - orbitals: ~1 is the 2s-state of Oxygen; 6 2 and 6 correspond to the interaction of the 3dz2 3 orbitals of Fe with the 2pz - orbital of oxygen; C5 4 is mainly the 4s - state of Fe. The oxygen atom bond with the chain is described by three in-

60

dependent interactions, the 2px - orbital of oxygen with the 3dxzband of the chain; the 2py - orbital of 0 with the 3dy z- band of the chain; and the 2s- and 2pz - orbitals of oxygen with the three bands of the chain presented in Fig. 14. The ~x - states of the FeO formed by the 2px - orbital of the oxygen and the 3dxz - orbital of the Fe (the same will be valid for the 2py- 3dyz - interaction) are separated by an energy gap ('" 5eV) which suggests rather strong Fe-O binding. At the same time, the 3dxz - band width is '" 2e V which corresponds to weak binding between the metal atoms due to these orbitals (i.e. overlap poorly). Thus, if the system comprising "2Px-orbitals of the o and the n3dxz-orbitals of the chain" are divided into two subsystems, Fe - 0 and (n-1) orbitals of the chain, they will be weakly bound. Hence, the energy spectrum of the whole system will comprise of two Jt x-levels of the Fe - 0 with the 3dxz -band of the Fe n_ chain between them. Since the ffx-levels of the FeO are far 1 from the band boundaries, the corresponding full states of the whole system will be localized; i.e. the atomic orbital coefficients of their wave functions will be at a maximum at the 2Px-orbital of the 0 and the 3dxz - orbital of the Fe (bound to the oxygen) and will then diminish exponentially in moving away from this pair of atoms. The results of the calculation (Fig. 16) agree with the scheme described. The energies of the levels of FeO are given in parenthesis. One can conclude, that from the point of view of the 2p x - 3dXz (2Py - 3dy z ) bonds, the adsorption of an oxygen atom leads to the formation of a quasimolecule of FeO only slightly disturbed by the other atoms of the chain. Comparison of Figs 14 and 15 shows that the 6 1 and 6 levels 3 of FeO (out of the permitted states of the chain) are converted into local states of the whole system and concentrated mainly on the corresponding orbitals of the oxygen atom and of the Fe-atom directly bound to the O. Another situation should be expected with (J 2 and 6'4 states; their energy is within the permitted bands, so they are converted into delocalized states within the band which are poorly bound to the FeO sUbsystem but cover the rest of the Fe-atoms. The results of the calculation agree with the scheme described (Fig. 17). There is a pair of states distinctly localized on the orbitals of FeO, their levels being close to the 6 1 and 53 states of FeO. In the region of "" 18.9 eV, there are

61

three or four states with very similar energies (they are shown boldly in Fig. 17); their wave functions are localized on the Fe - 0 to a greater extent than those of the other levels of the band. It can be considered as a single quasi-localized state describing the fragment Fe - O. Thus, most of the states determining the binding of the O-chain are localized ones corresponding mainly to Fe - O. Examining these states, one can easily examine the effects of the interaction of adsorbed species.

£=9.0 - - - - - [==/1.78

£=/5.1/

Fig. 17 - Scheme of the 3d 2z

3d 2

x -y

2-' 4s - and 2pz- sta-

tes of the chain with one adsorbed atom.

- - - - - [=J4.3JeV

Fig. 18 - Scheme of the 3dxz_

~

,==/4./3 £=/4.72

_C=I6.47

2Px(3dyz - 2py) states of the chain with two adsorbed atoms. (Changes in the coefficients of the wave functions for the local states are also represented).

62

Fig. 19 - Scheme of the 3d 2-' 3d 2 2 ' 4s z x - Y 2s - and 2pz- states of the chain with two adsorbed atoms.

t=!!J.48

==== C=.14. 24

e=J4.43eV

At large distances between the chemisorbed O-atoms, the above states do not affect each other. In mutually-approaching chemisorbed atoms, the fragments begin to interact through the chain and this causes a splitting into degenerate states; the wave functions of the split states begin to cover both fragments. Table 4 presents the splitting of the localized levels for various distances between the FeO-fragments. Figs 18 and 19 show the distribution of the one-electron levels and the wave functions of the localized states for a minimal distance between the FeO-fragments. In this case, the interaction between the two O-atoms is seen to be so significant that one can consider the combined complex to be Fe 202• TABLE 4

SplitUng of the Lower (LIef!, L1 e",) and upper (Je;.', J Cor ) Localized States of the 3dxz, 2px and 3d 2' 3d 2 2' z x - y and 2pz Levels (eV) at Various Distances, R, Between the Adsorbed Atoms (a is the Minimum Distance Between Fe-atoms)

1a

2a

4a

6a

t1 c 1

0.70

0.15

0

0

Llc 2

0.09 0.10

0 0

0

Lle 3

0.59 1.81

Llc 4

0.19

0.03

0

0

0

63

40

o

'

}

~ "

'0-

I

I

I

Cr

Mn

Fe

--'0

Fig. 20 - Calculated metaloxygen energies for the complexes Me~O"Me (1), ge(2), 0-0 0-0 M~ M~ (3), ~e/ (4) /46/

4< I

Co

,

Ni

Nevertheless the interaction between chem:'.sorbed atoms influences the position of one-electron levels of the system "metalchemisorbed atoms" (especially, the position of the localized levels appearing in chemisorption), the energy of interaction of the chemisorbed atoms (within the accuracy of the method used) is close to zero ~ven with chemisorption on neighbouring metal atoms. This result is in agreement with the data on the independence of qs on & for the system 02/Fe (Fig. 10) and suggests that if qs did depend on &, this fact should be accounted for by nonuniformity of the surface, different forms of adsorption, etc., but not by mutual interaction of the adsorbed species. Using the valence bond method, the author of reference /46/ showed that the oxygen-metal bond energies, c', are much higher for atomic oxygen than for molecular oxygen (Fig. 20). For each , . metal, the t values decrease a.n the order:

o

I

I

Me>Me

I

Me

""/

0--0

o-~o

>

Me

This is in accordance with the treatment discussed above for dependences of qs on 8 for the 02/Ag system /51/. It is interesting that the calculated c' values for different oxygen species change in parallel over the range of transition metals (Fig. 20) and are correlated with changes in experimental q~ values. I also calculated oxygen-metal bond energies for clusters containing several atoms of the metal (V, Cr, Fe and Co) which is a model of a real surface. It was shown that the most preferable configurations are those in which the oxygen was bonded

64

with one of the metal atoms of the first surface layer. The influence of sulfur additives on the chemisorption of oxygen on silver was explained by the fact that sulfur causes changes in the ratio of different forms of chemisorbed oxygen which results in changes in values of qs at various values of 9 /51/. Kinetic Methods of Estimating the Stability of Surface Species These methods are based on the validity of the Bronsted-Temkin relation in respect to the processes of the type (111.2) - (111.4). This allows one to judge the changes in q using kinetic parameters (the reaction rates or activation energies). Thus, by studying the isotopic heteroexchange between O2 and adsorbed oxygen

one can evaluate changes in oxygen-catalyst bond energies. The rate of the above process is often determined by the rate of oxygen desorption in which oxygen-catalyst bonds are broken /6/. The exchange between O2 and the oxygen adsorbed on Pt, Pd, Rh, Ir or Ni proceeds rapidly even at low temperatures (z OOC). The exchange is accompanied by irreversible oxygen adsorption, the rate of exchange decreasing progressively /62/. Oxygen adsorbed on platinum metals is energetically nonuniform; it occurs only to small coverages (1-5% of a monolayer) and is highly reactive due to low oxygen-catalyst bond energies (see Table 2). On Ni films at - 78°, only 15% of the adsorbed oxygen takes part in the exchange. Significant oxygen sorption accompanies the exchange and NiO is formed at 250°C. Bulk oxidation is also observed with Pd /62/. The heteroexchange on Ag starts at higher temperatures, proceeding rapidly at 200°_250°C. The exchange with oxygen on Au is observed only at temperatures exceeding 500°C /62/. Selenium additives decrease oxygen-silver bond energies (the decrease is 13-18 kcal mol- 1) as shown by exchange data /63/. The rates or activation energies of the surface reduction processes (111.3) and (111.4) can also serve as a measure of oxygencatalyst bond strengths /64/.

65

The above processes on Cu, Ag and Au are not obscured by HZ or CO adsorption. The interaction of these gases with oxygen, adsorbed on silver, at 170 0-280oC obeys a simple one-step mechanism expressed by Eqs (III.J)and(III.4). The reduction rate increased with 8. The activation entropy in the reduction with H2 (-35 cal deg- 1 mol- 1) and CO (-44 cal deg- 1 mol- 1) remains constant at different values of 8 which indicates the determining role of the activation energies. Since the heat effect in (111.4) is ~10 kcal(mol)-1 higher than that in (III.J), one should expect lower values of ECO in comparison with the values of EH at equal values of 8 • The 2

experimental data agree with the theory. Similar results were obtained in the reduction of surface oxides formed on gold /7J/. Metal Alloys Metal alloys can alsO catalyze reactions involving 02. In these alloys, complex phases are formed, the most wide-spread of which are solid solutions and intermetallic compounds /1-J/. Metals which have similar electronic structure of atoms and metallic radii, as well as similar crystal lattices form solid substitutional solutions. When atomic diameters differ in no more than 8-10%, full mutual solubility is observed; with higher differences, restricted solubility takes place. Metal atoms in a solid solution are distributed randomly in the crystal lattice according to statistics. In contrast to this, the alternation of metal atoms in intermetallic compounds (such as the phases CUJAU and CuAu) is strictly ordered. The probability of forming chemical compounds is higher when the constituent metal elements are situated far from each other in the Mendeleev Table. There are two types of intermetallic compounds: daltonides (constant chemical composition) and bertholides (variable chemical composition). The electronic structure of metal alloys is similar to that of metals. For catalytic reactions involving O2, the alloys based on noble metals are of the greatest interest. The alloys of the Pt-Ag type are interesting since they are formed by a metal atom with partially filled d-shell (Pt) and an atom with a completely filled d-shell (Ag). By changing the Pt/Ag ratio in Pt-Ag solid solutions, one can vary the concentration of holes in the d-band to elucidate the significance of the electronic factor.

66

Values of the metallic radii of Pd and Pt are rather close to those of the metals of the copper subgroup. In accordance with this, Pd and Pt give solid solutions with CU, Ag and Au. Palladium forms continuous solid solutions with the latter metals (full mut ual solubility). Their crystal structure is jsomorphous with Pd /2,74/. Platinum gives continuous solid solution with CU, while with Ag and Au, restricted solutions are formed. Copper gives restricted solutions with silver and continuous solutions with gold

/2, 71/. Using the method of surface equilibria, the authors of reference /56/ determined the bond energies of oxygen with Pd-Ag alloys (Fig. 21). One can see that addition of Ag to Pd leads to an increase in qs.

o

50

jl,. _

_

- 0 ....

P---r5--o-

, JO

40 I

Fig. 21 - Oxygen bond energies with the surface of Pd-Ag alloys /56/

I

20

40

60 Ag/wt. %

According to reference /62/, the introduction of Au into Ag (up to 50-60 at. ~ Au) enhances the rate of oxygen heteroexchange. Further addition of Au results in diminishing exchange rates (Fig. 22).

......0

e'\;

!I

~

to g

\ \ \ \

~

8

o

Fig. 22 - Rates of isotopic exchange of oxygen (molecules cm- 2 s-1) on Ag-Au alloys at 250 oC: 0 homomolecular exchange (PO = 0.08 Torr); • - ini2

tial heteroexchange (Po 20

40

0.1 Torr) /62/.

2

=

67

Metal Carbides Recently, beginning from the studies of N.I. Il'chenko /75-79/, the carbides of transition metals were used as catalysts of the gas-phase reactions involving molecular oxygen. The interaction of transition metals with such nonmetals as C, B, N, P and Si leads to solution phases. The latter keep the crystal lattices of the metals, small atoms of the nonmetals being distributed in intermetallic vacancies. These phases retain metallic properties: simple lattices, high electronic conductivity and high melting points. They have high thermal and chemical stability. The structure and properties of transition metal carbides /80, 81/ are of special interest. Solution phases are formed by the transition metals of IV-VI groups (TiC, ZrC, lifC, VC, NbC, TaC, cr 2, M0 2C, WC and W2C). In 3c the formation of triese phases, some of the valence electrons of the carbon are transferred to the d-levels of the metal so that the interatomic bonds in the carbides also include the electrons of the C. The melting points of the carbides exceed those of the corresponding metals. The carbides of the metals of group VIII (Fe C0 and Ni 3C 3C) 3C, and also Mn C form structures which are close to those discussed 3 above, but they are more complex. In the rhomboidal lattice of Fe the Fe atoms give trihedral prisms, the centers of which 3C, contain separate carbon atoms. The carbides of Co and Ni are unstable. The carbides of Fe, Co, Ni and Mn have lower melting points in comparison with the carbides of the metals of groups IV to VI. It has been shown /78/ that the bond energies of oxygen with the surface of the carbides is proportional to the heats of formation of the corresponding metal oxides, -~H~98. Fig. 23 demonstrates the linear relationship between the energies of the metaloxygen and metal-carbon bonds /78/. Carbon Catalysts Various types of activated charcoals can catalyze the reactions involving 02 (see, for example, /82/). Activated charcoals are porous sorbents with highly developed surfaces. Their skeleton consists of friable and randomly packed

68

-A Hfg8 (ltfBC)/IrC(J[(g-atomC)-1

50

40

Fig. 23 - The dependence of the heats of formation of the carbides upon those of metal oxides (Thermochemical data of references 17/and/8f).

JO 20 10

o -10 50

100 125 l -,1Hk8 (MBof!;-atOr

75

sets containing nets of six-membered carbon rings which are less ordered than in graphite. The nets are covalently bound by carbon radicals, hydrogen and oxygen /83/. Surfaces of specially prepared oxidized charcoals contain various functional groups (carboxyl, phenol, etc.) which are able to interact with reagents /84/. The bond energy of. oxygen with coal surfaces is likely to be close to that with transition metals (Co, Ni, etc.). The heat of oxygen chemisorption on graphite is 49 kcal (g-atom 0)-1 /16/. Oxide Catalysts Simple Metal Oxides Structure and Physicochemical Properties Crystal lattices of metal oxides contain ions of metals, Me n + , and oxygen, 02-, but the chemical bond between the ions is covalent to a considerable extent. The contribution of ionic bond is the highest when the electronegativity of Me n+ is low /85/. The types of crystal lattices of the oxides vary considerably and are determined by the cation nature, its degree of oxidation and ionic radius. Simple cubic lattices of the NaCl type, hexagonal lattices of the ~ - A1 20 type, tetragonal lattices of the 3 Ti0 2 type and monoclinic lattices of the Mo0 2 type are common for

69

solid oxides /86/. In the formation of a crystalline oxide from the constituent ions, splitting of the electronic levels takes place and this leads to definite bands corresponding to the levels in the isolated ions. In ideal crystals of the ZnO type, such bands are either completely filled or empty. The upper boundary of the highest filled band is separated from the lower boundary of nearest empty band (the conduction band) by a forbidden band whose width, U, is the energy required for electron transfer from the filled (valence) band into the conduction band. The number of electrons transferring into the conduction b-and increases with temperature (/V exp (- U/kT). The electrical conductivity of metal oxides is essentially less than that of metals so that the oxides are semiconductors. The above discussion refers to intrinsic semiconductors. In real crystals, there are many defects which generate discrete local electron levels between the valence and conduction bands. If the levels are occupied by electrons which are able to transfer to the conduction band at elevated temperature, one will have n-type semiconductors (i.e. electronic semiconductors). If the local levels are vacant but can be filled (by heat excitation) by electrons from the lower valence band, positive holes appear in the valence band which causes hole conductivity (hole, or p-type semiconduc tors). The work function, ~,in semiconductors is equal to the energy of electron abstraction from the Fermi level into vacuum (i.e. to infinity). The Fermi level is situated half-way between the local level and lower boundary of the conduction band (n-type semiconductor) or between the local level and upper boundary of the valence band (a p-type semiconductor). At elevated temperatures, real crystals manifest both intrinsic and defect electrical conductivity. When the latter is absent and U is very high, the crystal is an insulator. In n-type semiconductors, donor electron levels may appear as a result of anion vacancies (e.g. 0 2- are partially absent). To achieve compensation of charges, some cations gain lower degree of oxidation. For instance, in Y20S' y 4+ ions appear /87/. Another possibility is the presence of an above stoichiometric proportion of electron-donating metal atoms, as in the case of Zn atoms in ZnO. When there are more than the stoichiometric number of oxygen

70

ions, some of the cations gain a higher degree of oxidation (e.g. Ni 3+ in the case of NiO) which gives electron-acceptor levels and causes hole-conductivity. It is noteworthy that the simple band theory is not always valid for a description of the oxides of transition metals. For example, the oxides with simple cubic lattice, MnO, FeO, CoO and NiO, display properties which are determined mainly by the 3d-electrons of the cations. Each cation is surrounded by an octahedral field of oxygen anions in which, according to the crystal field theory, 11102 0 Mr;O 0

src;

o[aO oBeO 0ZrOz

ALzOJ 0 ~Oz 0 ThO 7il1l0 T/O

)(/0 o [O~

2000

o

.

reo

oCeOz

/000

.50

/00

Fig. 24 - The dependence of melting temperatures of metal oxides /86/ on their sublimation heats /31/.

(50 LjKc/I/·mot-/

the d-Ievels are split into t 2g (triply degenerate) and e g (doubly degenerate) levels. In the crystal, the above levels give bands; these are narrowed so that the t 2g - band is separated from the e g- band by a narrow gap (the forbidden zone). Semiconductivity will appear only in the case when the t 2g-band is completely filled while the eg-band is empty, for which each cation should have six electrons. This requirement is satisfied in only FeO. The semiconducting properties of the other oxides mentioned cannot be explained by the usual band theory. Some oxides (V20 V va, V60 1J' Ti 20 J, Nb0 2, etc.) display J, 20 4, metallic conductivity at low temperature. At elevated temperature, the electrical resistance increases suddenly; this is caused by lattice rearrangement and the appearance of semiconducting properties. The stability of metal oxides against melting (or sublimation) or dissociation is of great importance in catalysis. Many oxides of transition metals have rather high melting points which change in parallel to the sublimation heats (Fig. 24).

71

A measure of the ability of metal oxides to dissociate to form phases of lower oxides (or metals) is the change in the standard Gibbs energy for the process: (IlL33)

- 4 Gjge?crrt(f- at OF I 150

700

50

/00 _ !J;fg~Ctr{&_(JtlJ)-1

Fig. 25 - Standard free energy changes for reaction (III.33),kcal(g-atom)-1 of added oxygen, against the standard enthalpy changes for the transition: 1 - Ag-Ag20; 2 - cr20 3-cr0 2 ; 3 - Hg20-HgO; 4 - 000-C0 304; 5 - Pt-PtO; 6 - Cu20-CuO; 7 - BiO-Bi 20 8 - CU-Cu20; 3; 9 - MnO-Mn20 10 - FeO3; Fe 11 - Ti 304; 305-Ti0 2; 12 - Mn20 - Mn0 2 ; 13 - FeO3 Fe 20 14 - Cs-Cs 20; 15 3; 16 - Ga-Ga 20; Nb0 2-Nb20 5; 17 - Ga20-Ga20 18 - Cr3; Cr 20 19 - Mo0 2- Mo0 3; 3; 20 - Nb-NbO; 21 - Te-Te0 2; 22 - Sb-Sb 20 23 - Co-CoO; 3; 24 - Ni-NiO, 25 - Ge-GeO; 26 - Cd-CdO; 27 - Fe-FeO; 28 - Mo-Mo0 2; 29 - Al-A1 20 3; 30 - GeO-Ge0 2; 31 - SnO33 Sn0 2; 32 - Ce-Ce 20 3; Zn-ZnO; 34 - Ce20 3-Oe0 2; 35 - NbO-Nb0 2; 36 - MnMnO; 31 - Nb-NbO; 38 - TaTa 20 39 - Pu-PuO; 40 3; Np-Np02; 41 - Ti-TiO; 42 Ra-RaO; 43 - Hf-Hf0 2; 44 Sr-SrO; 45 - Mg-MgO; 46 Th-ThO; 41 - ThO-Th02; 48 - Ca-CaO.

72

Fig. 25 shows that there is a linear relationship between the values of L1 G~98 and i1 ~98 (per 1 g-ato m of added oxygen) and this means that there is approximate constancy of J S~98 values over the range of metal oxides. Similar relationships are observed at other temperatures /89/. Thus, the ability of the oxides to dissociation can be characterized by the values of j H~98/z. Fig. 26 presents the dependence of the standard enthalpies of formation of some oxides per 1 g-atom of Me on the n/m ratio. Slope of the curves is equal to j H0 298/z /90/. The latter is seen to decrease with increasing degree of cation oxidation which

Fig. 26 - The dependence of heats of MemO n formation on the n/m ratio /90/: 1 - Ti; 2 - V; 3 - W; 4 - Mo; 5 Cr; 6 - Mn; 7 - Fe; 8 - Co; 9 Bi; 10 - Ni; 11 - Cu; 12 - Rh; 13 - Pt; 14 - Ag.

indicates a greater ability to dissociate. The smallest slope is peculiar to higher oxides of Cr, Mn and Bi and also to the oxides of Ag (and Au); these oxides are really unstable under the conditions of catalysis. The most common effect of the reaction mixture over metal oxides in catalysis is the bulk reduction of the catalyst by the oxidizable molecules, R; this is dependent on the chemical nature of R, the temperature and the R/0 2 ratio, etc. /91/. The reducibility (in thermodynamic sense) is determined by the values of .1 GO for the processes:

73

This last equation is the difference between the equation of the process

o

ZIT(gflS)+l, :

='

zlTO (gas)

and that of process (111.33). If the same reductant in each case, (111.35) is constant. Hence, the l i ty (like dissociative ability) value for process (111.33), i.e.

one conSkuers the reduction by the L1 GO values for pro cess relative thermodynamic reducibiwill be determined by the LJ GO by the - L1 H~98/Z values (the

slopes of the curves in Fig. ?~). Comprehensive data on the thermodynamics of oxide dissociation are given in reference /89/. The Thermodynamic Characteristics of Surface Oxygen on Simple Oxides Measurements of the values of qs for metal oxides are complicated by many difficulties caused by the dependence of qs on 8 and, consequently, on pretreatment of the oxide. In calorimetric studies, an influence on the qs values of diffussion in the solid oxides has been found /92/. An example is given in Fig. 27 for NiO. In these experiments, a definite amount of oxygen was firstly removed from NiO by the reduction with CO. The 02 adsorption then took place and its heat was measured at 3270C /92/. As seen from Fig. 27, the deeper the preliminary reduction, the lower is the value of qs obtained (at fixed 8). The true values of qs decrease with growing 8. If the 02 adsorption took place on the most deeply reduced sample (curve 5), oxygen bulk diffusion of oxygen from the bulk to the surface had time to proceed during the experiment and this caused the true value of8 in the 02 adsorption ( e' ) to be higher than that calculated based on CO consumption in the preliminary reduction (e P ). The measured q values s 1 are therefore lower than those which would correspond to 8". At low degrees of preliminary reduction, the above diffusion effect is small ( 8' ~ 8" ) and the measured values of qs are close to the true ones (see Fig. 27). It is evident that the importance of the diffusion factor is determined by the mobility of oxygen in the oxide which depends on the chemical structure of the oxide and the temperature. The effect discussed is also essential when heats of the reactions (III.3) and (III.4) are measured. Insufficient rates of equilibra-

74

4

5

40

2 - 3.8%; 3 - 1.1%; 4 - 9.8%; 5 - 15.6% of a monolayer.

JO 2/

Fig. 21 - The effect of the depth of preliminary surface reduction (1-5) of NiO (B ° = 1 - S) on the heats of 02 adsorption at 321°C /92/: 1 - 1.6%;

L..-_ _--l-

5

.l.-._ _---I

/0

tion of the surface concentrations of oxygen may result in localization of the reduction on particularly active sites, leading to the formation of a lower oxide (or a metal) phase /92/. The phenomena described are similar to those observed in measuring adsorption heats on metals /21/. Systematic investigations of the values of qs for metal oxides using the temperature dependences of the eqUilibrium pressure of desorbed oxygen were made in references /93-95/. Fig. 28 shows

a5 1.0 1.5 2.0 80 / %

Fig. 28 - The dependence of values of qs' kcal(g-atom 0)-1, on the amounts of oxygen removed from the oxide (& 0' % of monolay4 - NiO; 5 - ZnO; 6 - V20 ; er): 1 - Cr 20 2 - CuD; 3 - Co 5 3; 30 4; 1 - Fe 20 ; 8 - Mn0 2 /93-95/ (the scales I to IV correspond to the 3 curves 1 to 4). the energetic nonuniformity of the surface oxygen. Similar results were obtained in reference /96/ for Mn0 2 and V20 • Initial values 5 of qs (q' s) referred to completely oxidized surfaces «(9 0 = 0) are related to the corresponding Po values (Fig. 29). This sug2

gests that there is an approximate constancy of the entropy factor in the 02 adsorption which is similar to bulk oxidation (Fig.

25).

75

Log/Po?11orr) 5 ZO(0204

o

[ilO

O

MO Crz0fe OJ

Ch20PIJOCdOo

t;

MfeA

117[1204

l4/j l::.

Mqf 8

UFBA

-5

-10 '--_ _----'!O

o

Fig. 29 - The dependence of the equilibrium pressure of desorbed oxygen (300 0C) on the initial heats of oxygen desorption /93-95/.

-'-_ _----=~ 20 ,]0

q;/!rC(lL(g-atr'

In reference /97/, Eq. (111.6) was also used for the calculation of qs at 8 = const. It was shown that in the case of V20 / 5 / A1 2 0 3' the values of qs are independent of e while for V205/3i 02' the values of qs decrease with growing 8. The values of q's for the rare earth metal oxides and for Mn0 2 were measured in reference /98/. The q values found by way of the method of adsorption-chemical s equilibrium /13,54,99/ are essentially higher than those of q~.This is caused by lower values of e in the experiments made by the above method. Systematic studies on the bond energies of surface oxygen using calorimetric measurements at 200 0-4000C were carried out in references /92, 100, 101/. Fig. 27 shows that the minimum value of qs (q's) for NiO at 327 0 C (the sample was prepared by the decomposition of the carbonate at 900 0 C ) was 22-23 kcal(g-atom)-1. At higher eo (lower 8), the values of qs increase, approaching the value for the bulk transformation of NiO into Ni (57 kcal(g-atom)-1). At lower temperatures (1500-2180 C ) , the initial energies of the desorption of O2 were considerably lower than the above value of q's which is explained by the changing form of oxygen chemisorption (the molecular form is prevailing). In the case of NiO prepared from the nitrate, the qs - 8 0 dependence (Fig. 30) is similar to that observed with the former sample. Lower values of q s for previously-reduced NiO are due to the influence of bulk diffusion of oxygen to the surface during the calorimetric measurements. Fig. 31 presents the qs - 80 dependence for NiO prepared from the 0C. carbonate at 600 In this case, q' = 11 and (q) = 54 kcal _ 1 s s max (g-atom) • Intervals in the experiments on oxygen removal by CO

76

(the times of the intervals are given in Fig. 31) result in decreasing values of qs' and this is caused by increasing B (i.e. diffusion of oxygen is esential). Fig. 31 also shows that there is an agreement between the values of qs obtained using Eq.(III. 6) (see Fig. 28) and those measured calorimetrically. A similar picture is observed with C0 /100/ (Fig. 32). In 304 this case,q's ~ 9 kcal(g-atom)-1 and the maximum value of qs is 35-40 kcal(g-atom)-1. r;s/KCtIl(g--tIt)-/

f

40

Fig. 30 - The dependence of the bond energy of the surface oxygen on the degree of surface reduction of NiO: 1 - qs values measured after removing ""3 and 6% of surface oxygen; 2 - the same after removing '" 13% of surface oxygen /92/ •

.10

5

10

00 / %

_-::;;.,;:;---~

50 40

fO 20 ':-_~_--':-_----'=----::-'=--:::--:'='"

20 00/% 0C

Fig. 31 - The dependence of the bond energy of the surface oxygen on the degree of surface reduction of NiO: 1 - qs values calculated from the heats of reduction with CO at 327°C; 2 - the heats of 02 adsorption at 3270 on a reduced surface /92/; 3 - values of qs obtained from the temperature dependence of the equilibrium pressure of 02 /93/.

For Mn0 2 at 200 and 8 0 <. 6%, qs = 18-19 kcal(g-atom)-1; qs is practically independent of 8 /92/. The values of qs for CuO at 300 0C and B 0 < 45% are also constant and equal to 34 kcal(g-atom)-1 /101/. Lower values were obtained using Eq. (111.6) (see Fig. 28). Similar relations are also found with Fe 20 /100/ (Fig. 33)~ 3 The important effect of oxygen diffusion is also observed. The

77

value of q's from calorimetric experiments is 20 kcal(g-atom)-1 and the maximum values of qs is 45 kcal(g-atom)-1. For other samples /93/, it was found that qs = 55-60 kcal(g-atom)-1 at 60 ~ 1070. The latter is close to the heat of the reaction:

Measurements of heats of 02 chemisorption on Mo0

J ~ ~

eo ••

c·

2

30 20 10

1 O~----'----_---L..._-_----I-

5

80 / %

10

3

at 400°C

Fig. 32 - The dependence of the bond energy of the surface oxygen on the degree of surface reduction of C0 1 - values of qs cal304: culated from heats of reduction with CO at 356°0; 2 - the heats of adsorption of O2 at 356°0 /100/; 3 values of qs calculated from the temperature dependence of Po /93/. 2

"" ~!

0-2

30

.-J

0-4

o

10

resulted in a large scatter of the data for qs' the mean value being equal~34 kcal(g-atom)-1 /92/. On Mo0 at 25°0, the 3/Al 20 3 values of qs fall from 30 to 15 kcal(g-atom)-1, the low bond ener-

78

gies being attributed to a molecular form of 02 adsorption /102/. Attempts have been made to standardize the q values for different s oxides. In references /93-95/, the initial heats of 02 desorption, , qs ' are proposed as a measure of the bond energy of the surface oxygen. These values characterize the completely oxidized surfaces (8 0 = 0) and refer to the most labile oxygen (see Fig. 28). The heats of oxide dissociation without the formation of new phases (calculated using thermochemical cy~les) were given in reference /103/ (see Table 5). These values of qs refer to the oxides which do not contain loosely bound extra-stoichiometric oxygen; they are therefore higher than the values of q~ • The values of qs for Mn0 2 and CuO are close to those obtained calorimetrically /93/ while the values of qs for the oxides of iron and vanadium approach the bond energies of surface oxygen measured using the method of surface reactions. TABLE 5

The Values of qs' kca1(g-atom 0)-1, for Metal Oxides

Oxide

qs

Oxide

NiO

13

Fe 20

Mn° 2

17

CdO

C0

21

V20 S Sn0 2

304 CuO Cr 20

3

34

3

Oxide

qs U0

qs

Oxide

qs

83

Mo0

3

91

61

3(U30 8) ZnO

83

Ti0 2

93

67

Ce02

85

zr0 2

96

70

W0

88

Th02

123

56

3

35*

*)Interpretation from the qs - q~

plot (Fig. 34).

The values of q s and q's are correlated (Fig. 34) (This re1ation was also noted in reference /104/). The observed parallelism between the energy of oxygen removal from the oxide (without phase changes) and the energy of desorption of weakly bound adsorbed oxygen is likely to be due to a similar dependence of the both values upon the common intrinsic parameter, C, the metal-oxygen bond energy, which is determined by the electronic structure of the cation.

79

Fig. 34 - The relationship between q s /103/ and q's /93-95/.

70 00 25

o

5

15

20

25 JO IJ§j/(Cfl{(7-flt;-1

Estimations of surface oxygen bond strengths can be expanded by using kinetic (relative) characteristics such as rates or activation energies of the isotopic heteroexchange reaction between 02 and the surface oxygen of metal oxides /105/ or those of the surface reduction reaction /64, 106/. It was found /93-95/ that the activation energy, Eo' of the former process for two atoms of surface oxygen is close to the initial desorption heats, q' (per s 1 g-mol of 2 ) :

°

(III. 36) If the surface oxygen is energetically nonuniform, the values of Eo refer to initial period of the heteroexchange reaction. Relation (III.36) allows one to obtain new values of q' ; for example, Eo for W0 is 51 kcal mol -1 /101/ and q's ~ 26 kc~l(g-atom)-1. 3 The activation energies of the heteroexchange reaction with CO 2(E CO ) can also serve as a measure of the bond strength of the 2

surface oxygen /108/ (see Fig. 35).

4:1lz. /lfCfll ·morl

Fig. 35 - Correlation between the activation energies of the heteroexchange reaction with 02(E o) and with CO 2(E CO ):

20

2

If}

o

20

40

50

Eo/!rcfll-morl

1,4,10 - V20 5 ; 2 - MoO); 3 - PbO; 5 Ce02; 6 - 3Ce02'V205; 7 - Fe 20 3; 8 cr20); 9 - CuO /108/.

80

2+10£T/H

!.O

'Z

Fig. 36 - The dependence of the initial specific rates of the surface reduction by H2 at SOOoC (I) and the 1 values (II) upon the TH 2 equilibrium pressure of desorbed oxygen: 1 - V20 ; 2 - V20 ( 80%) S S 3 - V20 ( 70%) - Mo0 Mo0 S 3(20%); 3 (30%); 4 - V20 ( SO%) - Mo0 3(SO%) 5 /109/.

40' oJ

0.5 2

j

o

1..J2

J

if

I..JO 1.28

/.0

1.5

The initial rates, r H ' of the reduction of the oxide with H2 2 (or the temperatures, TH ' at which equal initial rates of reduction are attained) were ~roposed /64, 106/ as being characteristics related to the bond energy of the surface oxygen. According to reference /103/, the values of T 1 decrease with increasing values of qs. Fig. 36 shows that the 2values of r H (at constant

H

2

temperature and PH ) correlate with the equilibrium pressures of 2 . desorbed oxygen. A similar correlation is observed for T 1• The

H2

corresponding activation energies of reduction, EH ' can also be used for the purpose discussed /110/. The depende~e of EH on 2

(Fig. 37) is similar to that of qs on 8 0 (Fig. 28). The above kinetic characteristics can be used only for a series of pure oxides. Small additions of other substances can influence the reduction rates (e.g. Pt in the reduction of V20 with H 2 S /111/) so that the r H values will depend not only upon q but (J 0

24f 20

°

2

.-

•

15- "

I

I

JO!

~!-) ~

14 o!t°

I

., ;;

I

!

"

zoe: · ~ .. ''5

II

2

If

J

5 7

.-.

12 !

I

I

"

J

2

, ,

1

/8F

I

!

2

_--

:~ r~~

t

I

,4 2

I

J

t

.JOt ,or~ ..x;8

4

ittl . . . . .·__. - -J

. . . lol,

-.--.-.1

I

I

J

°

I

~ 22 /

I

2

",.",0 2

.... 2'5

s

/ 22{ .- 18 ~;;-';.-.=---6

22 • 14,

J2F I

~~(~!

I

ae

I

I

I

I

I

I

04 0,5 o/J .. _o_o-q I

I

I

! !'

I

0.2 0.4 05 8g /%

Fig. 37 - The dependence of the activation energies of the reduction of oxides with H2 upon the amounts of surface oxygen removed, 8 0 / 110!: 1 - Fe 2 0 ; 2 - V20 ; 3 3 S co 0 ; 4 - Mn0 S - Cr 0 ; 3 4

2;

2 3

6 - NiO; 7 - CuO; 8 - ZnO; 9 - Ti02•

81

also upon the properties of the additions. A comparison of the qs values with various physicochemical characteristics of the metal oxides (like metals) is of definite interest. First of all, there is some correlation between qs and the heats of bulk oxidation of the oxides (Fig. 38). Many oxides t1H o -..-.l28/ /(C£l/(r;-£lt}-1

z

/50

ZrOzo

o ThOz

/00

Fig. 38 - Heats of bulk oxidation of metal oxides (see Eq.(III.33)) plotted against qs

50

25

50

75 100 ~sj;rc£l/&-tlt;-I

/25

(e.g. those of Cr, Mo, W, U and V) do not obey the correlation. In general case, therefore, the bulk thermochemical properties such as - Ll H~98/z cannot characterize the bond strength of the surface oxygen. The same applies to the standard heats of formation of the oxide per 1 g-atom 0, - Ll~98/n 1113, 114/, although there is some correlation between this value and qs (Fig. 39). Since the metal-oxygen bond energy affects the band structure of the metal oxide, one should expect that there might be a correlation with the width of the forbidden band, U. This is actually observed 1351 (see Fig. 40). This correlation should result in a correlation between qs and the colour of oxides, since, for semiconductors with simple band structures, the energy of the first optical transition coincides with U. In fact, the oxides with low values of qs are deeply coloured (Co Mn0 2' NiO, CuO), the oxi304, des with moderate values of qs are coloured less intensively (Fe 20 CdO, V20 and the oxides with high values of qs are only 3, 5) slightly coloured (U0 ZnO, Ce02, W0 Ti0 2). This rule is rough 3' 3, and there are definite exceptions (e.g. V20 U Sn0 2). In ad4, 30 8' dition to common intrinsic properties which determine both the absorption spectra of oxides and their qs values, there exist other factors which influence each of the above characteristics in a dif-

82

Fig. 39 - Heats of the formation of metal oxides plotted against qs·

o L.-

.l-

::':"::-_

ferent way. There is also some qualitative correlation between qs and the type of conductivity of the metal oxides. Oxides with low values of qs are predominantly p-type semiconductors; the oxides with very strongly bound oxygen are insulators; the oxides with intermediate values of qs are n-type semiconductors. fJ/eJi 4

S!7O

0

CdO

2

oNtO

feZOJ 0

Cti°o

o

MoO

OJ

'J

0 WO" tiD,!

Jlz°5

o

!jS/j(CU~

100 80 80

/0

~

~~

<, ,

.~L:>~ 2

~

5

40

COf04

/

~.J~

TiOz ZnOO

'2

o

20

MllOz

o

/0 20 JO 40 50 80 70 80 gO /00

!Js/Kcatfff-atf/ Fig. 40 - The width of the forbidden band, plotted against qs

Ti

v

Cr

Mn

Fig. 41 - Calculated values of the energies of the 0-MeO(1) and HN-MeO (2) bonds and the experimental values of qs (3) /118/.

According to reference /115/, the oxides with high values bf q s have the most stable configurations in the d-shell of the cation,

83

i.e. dO(Ti0 2, V20 d S(Fe 20 ) and d 10(ZnO); less stable confiJ S)' gurations are peculiar to the oxides with lower qs values, i.e. d J (Mn0 ) ' d 6, d 7(CO d8(NiO) and d 9(CuO). In addition, one 2 J0 4), should note that the unstable configurations d 8 and d 7 are also C 2+ an d N,2+, t 0 Pt 2- , Pd 2+ , Rh2+ , 0 t 0 ca t'~ons , pecuI ~ar ~ ,~.e. 0f the metals on which the heats of oxygen adsorption are low. In reference /116/, the simple crystal field theory was used to calculate changes in the contribution of the energy of the crystal field to the energy of oxygen abstraction from the cation in octahedral and tetrahedral coordinations. These contributions were shown to be predominant in determining the bond energy of the reactive oxygen in a range of metal oxides. Some attempts were made to estimate surface bond energies of the oxygen theoretically. One of these attempts is based on a purely ionic model /117/. The latter is evidently insufficient, since the Me-O bond is covalent to a great extent. Relation (111.17) suggests that for the estimation of changes of qs over the range of metal oxides, one can calculate energies of dissociation of simple complexes MemO n• Such calculations were made /106/ using the expanded Huckel method. Fig. 41 shows these values to change in parallel with qs. A more complex semiempirical scheme of calculation, where the oxide is considered as an aggregate of interacting Me-Q-bonds, was proposed in reference /119/. Bond Energies of Various Molecules with the Surfaces of Simple Oxides In contrast to bond energies of oxygen to surfaces, the corresponding data for other atoms attached to the surface are scarce. Bond energies obtained using the multiplet theory /12/ are hardly valid for oxidation catalysis to which the theory is not applicable. A variety of forms of adsorbed molecules makes the problem very difficult. For example, in the adsorption of propylene on metal oxides, the rupture of a C-H-bond, leading to an allylic complex, or that of a C-C-bond, resulting in carbonate-carboxylate complexes, are both possible. Besides these, $-complexes of the adsorbed olefin as well as weakly adsorbed (slightly polarized) molecules are also probable. Heats of adsorption measured by chro-

84

matography using Eq.(III.5) are likely to refer to the last two sorts of species. For example, in the adsorption of propylene on Bi 20 at _8 0 _ + 31°C, the heat of adsorption, q, is equal to 3 6.4 kcal mol -1, while on Mo0 (0-65 0C) q is 3.3 kcal mol -1 /53/. 3 The q values at 40 0C for different olefins on C0 vary between 304 6 - 14 kcal mol -1/143/. Similar difficulties appear in the case of more simple molecules. For instance, CO can be adsorbed weakly and reversibly or it can interact strongly and irreversibly with the surface to give carbonate species /16/. The latter are also formed in the adsorption of CO 2• The heat of adsorption of CO 2 on MuO is 20 kcal mol -1 /16/ which is close to heat change for the reaction MuO + CO 2 = = MnC0 According to reference /92/, qco for NiO (with 0.1 at.% 3• 2 1i) is equal to 25 kcal mol -1. Molecular hydrogen, like CO, can be adsorbed in two main forms. The weakly bound one prevails at low temperatures while strongly attached species are predominant at higher temperatures. The first adsorption proceeds fast and reversibly in contrast to the second one. In the second case, surface OH-- groups are formed and the metal cation is reduced, for example: H + 02~4+02-= OH-Mn2+OH-. 2 The desorption of such surface species leads to the liberation of water but not ~ (surface hydroxides are dehydrated); like-wise the decomposition of surface carbonates results in the desorption of CO 2 but not of CO /16/. The chemisorption of S02 on metal oxides is likely to result in the formation of sulfite or sulfate on the surface: S02 + (0) = (S03) or S02 + 2(0) = (S04). The initial heat of the adsorption of S02 on CuO is equal to 26 kcal mol -1; the value of qso gradually 2

attains 31 kcal mol -1 /16/. The latter is close to the heat change for the formation of copper sulfate from S02 and 2(0). Discrimination between heat changes referring to definite forms of adsorption is a rather difficult experimental problem which is complicated by the reduction of the metal oxide at the elevated temperatures at which the catalytic process takes place. This is why quantum-chemical calculations of the bond energies of the oxides with various surface species are of considerable importance (see, for example, reference /41/). It should be noted that a kinetic comparison of the bond strength of various molecules with oxides can be obtained from the relative rates of surface reduction of the oxide by different sub-

85

strates provided that the reduction mechanism is similar for all the reductants. Acid-Base Properties of Oxides In some cases, the surface intermediates formed during the course of oxidative catalysis involves salt-like surface complexes /121/. The latter appear, for instance, during the reaction of an oxide surface with organic acids, RCOOH, and are destroyed by interaction with OH-groups. The acid-base properties of various oxides can therefore be characterized /35/ by the heat change, qRCOOH' of the process:

~ ;l1e(IfCOO)Il($Olid)~:

Me (Oll)n (sotia )» /fCOOII(;dS) +-

(IIl.37) For various organic acids, the sequence of change in qRCOOH is the same and close to that of the dissociation energies of the carbonates (Fig. 42): on going through the sequence Cs+, Rb+, K+, Na+ to Ni 2+, Cu2+, Fe 2+, the acidity increases. To expand the f -qRCOOH/KcaL·mol-

90 80

70

60 50 40

[(/+~ 2 Co +Ml+

30 N/+ An+ /2.

20

cut-t

AJ

"/!l

o -f

1:J-2

v-s

z+

0-4

Cd 2+

fO ':- :':- : '- -=,"~'-:- '-:- f- -'- - J fO 20 JO 40 50

o

sa

70 80

-qcoz/lfcal-moZ-'

Fig. 42 - The dependence of the heats, qRCOOH' of reaction (111.37) on the heats, qco ' of dissocia2

tion of the carbonates for various acids: 1 - fo~mic acid, 2 - acetic acid, 3 oxalic acid, 4 - benzoic acid.

thermodynamic scale of acid-base properties, it is advisable to use also the heats, qHCI' of the process:

1.. Meet17 (solid )+1... Me (Ofl){soliti)=flCt({1dS)+-l. n 11 ';1' n ¥e.;z 011 (..fOlid)

(IIl.38)

86

r;HCI/Kcul-mo/-1

100 gO

Fig. 43 - The deperrience of the heats of process (lII.J8) on the dissociation energies of the carbonates.

80

70

50 50 40

JO

2+ C02+Pb Cd2+

co

Cfil+fet•

20 Ntb8 10

o

0

Z11 2+

o/r111 2+

MgZ+

/0 20 30 40 00 50 70 80 gO 100 -!JCOp/trCO'/-l11ot-1

which change in parallel to qco

2

and hence also qRCOOH (Fig. 43).

The above heats are also correlated with the Xi' /85/ of the metal cations (Fig. 44).

electronegativity,

Ii 0-/

/:; -2

15

jea]+ JOO

CrJ +

eZnZ

+

30

M!1Z+ fY;J+

'0

Ca Z+ 3r Z+

1.0

o

Na'

RbT CST

25

50

r 2+ GO

o CIJ.2+

100

f5J

o

0

#12+

75 /00 -flRCOOH /Kcat'l11ot- 1

Fig. 44 - Correlation of the heats of processes (111.37) with cation electronegativity ,z i : 1 - formates, 2 acetates.

0

t:::.

Mg 2+ L-_....L-_---L.._ _L--_.l-_.....J

10

20

20 10

0

JO 40 SO -f/HCL/XCtli'/110t-1

Fig. 45 - A comparison of increase in vibration frequencies of NH Ll J NH ' 0 ) and pyri3( 3 dine (L1 J J5H5i~' Ll on adsorption on various metal oxides /122/ with the heats of process (IUd8).

87

Using qHCl and ,xi' one can classify the acid-base properties of the cations /35/. The above parameters are also valid for the surface properties since there is a correlation between qHCl and the change in the vibration frequencies of Lewis bases (ammonia, pyridine) when they are adsorbed on metal oxides (Fig. 45). A kinetic characteristics of acid-base properties of the oxide surfaces can be obtained by measuring their catalytic activity in respect to such processes as n-butene isomerization /123/. In general, there is no correlation between the surface bond energy of oxygen adsorbed on the oxide and its acidity (Fig. 46).

MfiOZ

o

NiO

o

oJ!lOz o

FeZOJ

CoJ04

o

o CrZOl

1.5

ilz05

CI10

o

0

0

[dO

0

Fig. 46 - A comparison of the values of qs and the electronegativities of the cations.

MoOI WOJ

°°

VOl

lTlO°

TlO!

°°lr02

o

°

Th02

Ce02

!

o

25

50

75

100 /25 l q)/KCflL (7 -at

r

Complex and Modified Oxide Systems One can distinguish several types of complex oxide systems which are employed in catalytic reactions involving 02. Homophase complex oxide catalysts include those oxides which form new chemical compounds or solid solutions which are new chemical individuals (daltonides or berthollides). In other cases (heterophase catalysts), a complex catalyst includes the phases of two (or more) oxides or mixtures of their chemical compounds (solid solutions) with an excess of initial oxide. If small amounts of another oxide are added to a given metal oxide, the latter may be modified /124/ by the former. The modifying additive can chemically interact with the main component and this leads to a homophase system. However, the additive may also form a separate phase. The first type of complex oxide catalysts includes the spinels,

88

Mei+(Meii)20~-,

which can be considered to be a compound of MerO

and (Merr)203. Cobaltites, chromites, ferrites and manganites (Meii = C0 3+, cr 3+, Fe 3+ or Mn 3+ ) belong to this class. The metal ions mentioned form amphoteric oxides. Spinels of the Me r(Me r r)2 04 type are isomorphous compounds. Their unit cell contains 8 units of Mer(Merr)204. Anions of 0 2- form a compact cubic structure; the tetrahedral vacancies are occupied by 8 cations, while the octahedral ones are occupied by 16. The formula of the spinel will be

With a = 1, we have a normal spinel, while with a = 0, an inverse spinel is formed; with 0< a «,1 a mixed spinel is formed. The value of ~ is determined by ability of the cations to adopt tetrahedral or octahedral coordination which depends on the energies of stabilization by the crystal field /126/. Manganites and chromites are mainly normal spinels while ferrites arc reverse ones. Electron transfers may lead to ~alence exchange between Met and Me r r• For instance, iron cobaltite may contain C0 2+ and Fe 3 (but not C0 3+ and Fe 2+). The oxygen content may vary over a small range near to the stoichiometry corresponding to the given spinel phase /127/. The thermodynamic properties of spinels have been relatively rarely studied. The ferrites have been investigated in most detail /121/. Complex oxide systems formed by acidic oxides (e.g. Mo0 W0 3, 3, U0 V20 , P20 , As20 , Sb 20 and Se0 2) in combination with basic 5 3, 5 5 5 or amphoteric oxides (NiO, CoO, CuO, Fe 20 Bi 20 and Cr 20 are 3 3, 3) salts (molybdates, tungstates, uranates, vanadates, etc.). The structure of molybdates has been best studied due to their practical significance. For instance, in the Co-Mo-O system, two-phase regions have been distinguished, (C0 + CoMo0 and (CoMo0 + 4) 304 4 + Mo0 /128, 129/. Cobalt molybdate has two modifications diffe3) ring in the binding of the Mo0 6 6- octahedra /128/. rn the Fe 20 3 is formed. With - Mo0 system,normal iron molybdate, Fe 2(Mo0 4)3' 3 an excess of Fe 20 a two-phase system (Fe 20 + Fe 2(Mo0 exists, 3, 4)3) 3 while with an excess of Mo0 a two-phase system (Fe 2(Mo0 + 3, 4)3

89

+ M00 is detected /130/. The structure of the other trivalent 3) metal molybdates is similar to iron molybdate /131/. In the Bi - Mo0 system, three compounds, Bi 20 . 3Mo03 (at.. -phase), 3 3 20 3 Bi 2Mo0 (j9 -phase) and Bi 20 Mo0 ( -pnase) are found 3 3' 20 3• 3 /14/. The Sn0 2-Mo0 system may contain mixtures of three phases 3 (Mo0 tin molybdates and solid solutions of Mo0 in Sn0 2). The 3 3, Mo0 and an anSb - Mo0 system includes mixtures of Sb20 4, 3 20 4 3 timony molybdate /133/. Metal tungstates are similar to molybdates. Various vanadates have been also studied in detail /134/. Vanadium pentoxide combined with alkali metal salts (sulfates, etc.) give various vanadates and sulfovanadates as well as vanadyl - vanadates and "vanadium bronzes" /135/. Vanadyl - vanadate formation is accompanied by the liberation of 02 /136/. In the Bi 20 - Sb 20 system, the following phases are detected: 3 4 a solid solution of Sb 20 in Bi 203, a bismuth antimonate BiSb0 4 4, a solid solution of Sb 20 in BiSb0 and a mixture of Sb20 with 4, 4 4 the last solid solution. At various Sn:Sb ratios, solid solutions of Sb 3+ or Sb 3+ oxides in Sn0 2 or mixtures of Sb 20 with the above 4 solid solutions are formed /133/. In t he Fe-Sb-o system, three stable phases exist: Sb 20 Fe 20 and FeSb0 (at Fe:Sb = 1 : 1). 4, 4 3 At other Sb : Fe ratios, a heterophase system exists /137/. In the addition of Mo0 to V20 (up to 25-30 mol % of Mo0 a 3), 3 5 solid solution based on V20 is formed. Increasing the Mo0 con3 5 tent (up to 50 mol %) results in a heterophase system containing the above solid solution and a new chemical compound V20 Mo0 5• 3• At higher Mo0 concentrations, the system contains this compound 3 and MoO)/132, 138, 139/. The systems V20 S - P20 S' W0 3 - P20 S' V20 - we), etc. are likely to have similar structures /14/.

r

S

Compound oxide systems without mutual chemical interaction of the components (mechanical mixtures) are also used in oxidative catalysis. Catalysts of the hopkalite type, CuO + Mn0 2' or CuO + + Al 20 3 belong to this class of heterophase catalysts. However, one can assume that chemical interaction of the components occurs on the boundaries of the component phases. Homophase and heterophase systems can involve more than two oxides. For example, in bismuth molybdate, part of Bi 3+ can be sUbstituted by Fe 3+; hopkalite-type catalysts can be promoted with Ag20, etc. The modification of oxide catalysts can be also classified. For

90

instance, small additions of 1i 20 or Ga20 to NiO or ZnO give so3 lid solutions, while in other cases, such as the addition of Bi 20 J to CuO, neither solid solutions nor chemical compounds are formed /14/. Mixed systems are also possible. For example, Ag20 addition to Mn0 leads to silver manganite formation, and a separate Ag20 2 phase also exists simultaneously /140/. A new type of promoted oxide catalyst has been proposed by N.I.Il'chenko, V.A.Yuza and V.A. Roiter. These consist of metal oxides doped by small additions of platinum metals /141/. The promoter forms highly disperse crystals on the oxide surface /141, 142/ as seen from Fig.47 where an electron micrograph of V20 do5 ped with Pt is shown. Discrete Pt crystals which do not penetrate into the oxide lattice, are formed in contact with the V20 Simi5• lar catalysts can be prepared by tabletting mechanical mixtures of the oxide powder (e.g. V20 with Pt on an inert carrier (e.g. 5) BaS0 /141, 142/. As the promoters, not only metals but also me4) tal carbides or other metal-like compounds can be used /75/.

Fig.47 - Electron micrograph of a V20 catalyst with 0.5 wt. 5 % Pt. Magnification 3840 /141/.

Recently, other types of oxide catalysts have been employed widely. Solid solutions of catalytically active oxides (e.g.CuO, NiO, Cr 20 , CoO) in an oxide matrix (e.g. MgO) are of interest /144,145/. J With small concentrations of the active component, these form isolated sites which are similar to those involved in homogeneous catalysts. Isolated active centers also occur in metal-substituted zeoli t'es as well as in the catalysts in which active species are bound by chemical bonds and attached to an inert carrier. An example is given by vanadium oxide systems supported on silica gel or alumina obtained by chemical modification of the carrier surface /161/.

91

The Bond Stregth of Surface Oxygen in Complex Oxide Catalysts Vfuen constituent components of a complex oxide catalyst give a new chemical compound, the latter has a value of qs which is not the sum of the values of qs for the components. Fig. 29 presents the initial heats of oxygen desorption, q~, for spinels to illustrate this conclusion. The bond strengths of the surface oxygen for manganites and ferrites are given in reference /147/. The heat of oxygen chemisorption on Fe 2(Mo0 measured calori4)3 metrically is equal to ~41 kcal(g-atom 0)-1 for a wide range of surface coverages, B. The value of' qs for Fe 20 - Mo0 with 3 3 Mo : Fe = 1.75 is ""' 37 kcal(g-atom 0 )-1 /92/. Thus, complex Fe Mo - oxide catalysts are characterized by higher values of qs than Fe203 and by lower qs values than MoO). According to reference /92/, bulk oxygen in Fe - Mo - 0 is very mobile. In contrast to that, the rate of diffusion of bulk oxygen in cobalt molybdate (at 400°0) is low. The heat of adsorption of oxygen in the le-st case is 15 kcal(g-atom)-l at 6l o = .' f; { and 44 kcal(g-atom)-l at B 0 ::> 6010 /92/. Table 6 shows that the bond strength of surface oxygen increases with Mo content. TABLE 6

Isotopic Oxygen Heteroexchange with Co-Mo-Oxide Catalysts at Po = 10 Torr /147/ 2

Samples/ at.to Mo

° (Co3°4) 21.0 50.0 (CoMo0 4) 10.0 95.0 100 (Mo0 3)

Initial rates of exchange, r . 10- 11, at 600 0C (r in molecules cm-2 s -1) 100-125 350-400 500-550 600-650 600 525-600

2200 900 8.0 0.63 0.41 4.4

E/kcal (mol)-1

16 29 42 50 51

92

The data /148/ for the reducibility of the molybdates with ammonia are in agreement with the conclusions /92/ on the bulk diffusion of oxygen in the molybdates. The bismuth molybdates, Bi 20j .JMoO and Bi 20 are characterized by fast bulk diffusion J, J J of oxygen. The value of qs~J8-J9 kcal(g-atom 0)-1, irrespective of 8, in the both cases /92/. The addition of P 205 to bismuth molybdate diminishes oxygen mobility. In this case, the value of qs becomes dependent on 8 and decreases, with increasing 8, from 48 to JO kcal(g-atom 0)-1/ 92/. The heat of oxygen chemisorption on chromium molybdate and on Cr-Mo-O (Mo : Cr = J) is J7 kcal (g-atom 0)-1 at various values of 8. The above catalysts were partially reduced in the course of catalytic oxidation of methanol /92/. oMoO

TABLE 7

Isotopic Heteroexchange of Oxygen with Vanadium Oxide Systems Containing K 2S0 /154/ 4 System

Exchange rate

E/kcal mol -1

r 010 J /g m- 2 h-~

at

Po

440 0C 0.OJ2 0.20

4.70 2.40

= 40 Torr

2

480 0C 0.17 5.70 14.0 8.60

44 4J 31 J5

(0.32)

Reducibility with propylene /132/ suggests that the surface oxygen of various tungstates is bound more strongly than in the Bimolybdates. The value of qs for tin vanadate increases sharply with decreasing 8. The initial heat of 02 desorption from sn(VO is by J)2 8 kcal(g-atom 0)-1 lower than that from V205/96/. Thus, the value of qs for sn(V0 is higher than that for Cr 20 and lower than 3)2 3 that for Fe 20 (see Fig. 29). 3 The addition of Sn0 2 or Ti0 2 decreases the values of qs for

93

V /149/. A more complex effect is caused by the addition of 205 Mn0 to V20 5 /150/. 2 Potassium sulfate addition (up to 0.5 mol. %) to V20 5 results in a loosening of the V-O bond (Table 7). The effect of V-Q bond loosening by alkali metal sulfates increases in the range: Li2s04~Na2S04~K2S04"(Rb2S04,(Cs2S04/152/. Table 8 shows that other compounds of potassium cause a similar influence on the V-O bond strength. TABLE 8

Isotopic Heteroexchange of Oxygen with Doped Vanadium Oxide Systems /153/ Temperature of preheating in 02/oC

System

440 400 400 400 400 400 400 400 450 450 450 400 450

V20 S Li ZS0 V20 4 5 NaZS0 V20 4 5·0.1 V20 K 5·0.1 2S0 4 K V20 S,O.05 2S0 4 V20 Rb2S0 4 5·0.1 V20 Cs2S0 4 5 V20 KOH 5·0.1 V20 KOH 5·0.2 V20 7K2P0 4 5·0.06 V20 K 5·0.1 2S0 4 V20 0.1 K2S20 7 5· V20 K 5·0.1 2S20 7 00.1

00.1

Exchange rate 3, at 440 0C r ( . r J.n g m-2 s -1) ol0

0.03 (0.04) 0.1 1.42 0.72 1.75 5.26 (0.58) (0.68) 0.91 3.21 7.46 10.7

E/ kcal mol -1

46 46 45 42 42 41 35 33 33 33 38 38 38

The isotopic heteroexchange of oxygen with tungstates and uranates of sodium has been studied in references /154/and/155/. The variation of the bond energies of the surface oxygen for the Fe-Sb-Q sYstem is given in Fig. 48. The corresponding isotopic data are presented in Table 9. The heat of the chemisorption of oxygen on a Sn-Sb-Q catalyst increases from 35 to 49 kcal(g-atom 0)-1 with increasing 8 0 (~ 5%). At B '> 5% q :::: 49 kcal mol -1, irrespective of 8 / 92/. o

s

0

94

80 %

50 r'(07

(5

00 0

1

qs 35

10 30 5

It'ig. 48 - The values of qs for the Fe-Sb-O system (1) and the specific rates of reduction with 1-butene at 425 0C (2) for various degrees of reduction of the catalyst surface /156/.

25

0

20

20

40

00

80

100 80 /%

TABLE 9 Isotopic Oxygen Heteroexchange with Fe-Sb-O Catalysts at Po = 10 Torr /147/ 2

r.10:) / molecules cm-2 s -1 , at 550 0C

Sb/ at. %

° (pure

Fe 20 3) 25.4 39.0 54.3 63.9 66.0 66.7 91.9 95.0 100(Sb20 4)

325-345 340-407 450-530 548 547 549-588 530-598 551-579 552-579 538-600

E/ kcal mol -1

1100 320 310:t50 9.2 7.6 5.4 4.5 6.8 0.94

45 45 45 44 45

The values of q for V-Mo-o catalyst turned out to be ~34 -35 kcal(g-atom O)-~. High mobility of the bulk oxygen was also found /157/. According to references /109/and/139/, the bond strength of the surface oxygen, for systems containing up to 30-50 . mol. % Mo0 decrease with increasing Mo0 concentration. Further 3, 3 increase of the Mo0 content leads to increasing values of qs. 3

95

These results agree with the IRS-data /132/ but are in contrad.iction with references /158/and/159/. Thi8 problem requires further investigation. Small amounts of metal oxides can also change the bond strength of the surface oxygen. Thus, calorimetric studies have shown /92/ that the addition of 1i 20 to NiO (0.1 and 2 at. % 1i) decreases the bond energy of surface oxygen in NiO by 2 and 7 kcal(g-atom 0)-', respectively. A similar effect was observed when NiO was doped wi.t h Ag20 /146/. At the same time, heterophase additives like Pt do not influence the bond energy of surface oxygen in V20 and other metal oxides S /99/. In conclusion, one may note that the values of qs for spinels of the type MeI(MeII)204 are determined mainly by the properties of Me I I• For example, the values of q~ for Znco 20 is close to 4 that for co and the values of q~ for NiFe 20 MgFe 20 and 4 4, 304 CdFe20 are close to that for Fe 20 (see Fig. 29). Similarly, the: 4 3 bond strengths of the surface oxygen on molybdates, tungstates, vanadates, etc. is determined manily by the MoO), W0 and V20 5, 3 respectively. The value of qs for NiO doped with 1i 20 is close to that for NiO. These cqnsiderations can be useful as a first (and rough) approximation.

Metal-Substituted Zeolites Zeolites containing transition metal ions can be obtained by way of ionic exchange, for example, between NaY and Fe 2 + , Co2+ , 8 Ni 2 + , Cu2 + or Ag+ /160/. Isotopic heteroexchange of oxygen with zeolites of this type has been studied. The activation energy for NaY is equal to 45 kcal mol -1 while for the Ag+-zeolite it is equal to )8-39 kcal mol -1 and for the Cu2+-zeolite, it is 2) kcal mol -1/88/.

01

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