Bonding in Transition Metal Complexes

Bonding and elementary steps in catalysis

4.1 INTRODUCTION

4.2.2 Bonding in Transition Metal Complexes

4 ~ BONDING AND ELEMENTARY STEPS IN CATALYSIS 117 is proportional to the number of nearest-neighbour atoms) is small compared to the adsorption energy to an open surface (surface atoms have low co-ordination).

Note that the adsorption energy dependence on surface atom co-ordination relates to the weakening of the bond energies between the surface atoms (due to electron localization, only partially being restored by Eemb).

When intermediate surface bonding (~fd[3 = ~/z~) occurs the adatom bonding and antibonding orbitals are not any more separate, but then overlap as broaden- ed bands in the surface electron density regime (Fig. 4.3d).

In the weak interaction limit (Fig. 4.3e) the orbitals collapse. The localization energy cannot be overcome and the surface electron structure is only weakly disturbed.

The band width of the electron distribution on the surface atoms (the analogue of A) has now become very small and equal to:

__ 2 , f ~ , ~ '

A 2nl~'l = 2 "(4rnl i) (4.15a)

<< ~/-dl [3' [ (4.15b)

The covalent contribution to the bond energy has now become very much less than the free cluster value ~Fn 113' i. The bond energy decreases ~(~/~)-1 when the surface co-ordination of surface atoms increases.

118 4 ^-^- BONDING AND ELEMENTARY STEPS IN CATALYSIS

The relative energies of C O valence orbitals and Ni atomic orbitals are sketched in Fig. 4.4b. The Ni a t o m has eight occupied 3d a n d t w o occupied 4s orbitals.

Because of the difference in spatial extension, one can ignore h y b r i d i z a t i o n b e t w e e n d and the s, p electrons. As in tetrahedrally co-ordinated carbon atoms, the 4s and three 3p orbitals can be h y b r i d i z e d to four sp3 orbitals (Fig. 4.4a). This

CO CO

,ii

x " ~ C O

CO

Fig. 4.4a. Geometry of Ni(CO)4.

E

free atom

{* antibonding }~

//

:-' l + + -~A !antibonding

"I./ ... dx, d,z dzx~- i

J \ ... d:,2 d x' y2 ~J&- - . . . . ::::.":. ... - "- I , \ "~

- \ \!

9 ,... -... /

. . . . . . ..

':'\ ___. "bonding complex

2~"

5~

molecules CO

Fig. 4.4b. a-interaction in Ni(CO)4.

4 ~ BONDING AND ELEMENTARY STEPS IN CATALYSIS

...r -..

..." ...

E ..-

+ + ...

d x2_y2 dz2 ...

... +4_/'

antibonding

: 2/1:"

bonding

119

no 2~" interaction interacting

system no i nt er ac t ion with C O 2~ "orbital

Fig. 4.4c. ~-interaction in Ni(CO)4.

will push their energy upwards with respect to the 4s orbital of atomic Ni.

Therefore the two 4s electrons present in the free Ni atom will transfer to the empty 4d orbital and the Ni 4d electron system becomes dl0.

In the N i ( C O ) 4 complex the four empty Ni sp3 orbitals are directed towards the four doubly occupied 5(~ CO orbitals, with which they form four occupied bonding and four unoccupied antibonding orbitals, separated by bandgap para- meter z~. Because of the large overlap between the Ni sp3 orbitals and the CO 5r~

orbitals the resulting covalent bond will be strong and dominate the Ni-CO bond interaction. Because the CO 5r~ orbitals are doubly occupied and the Ni sp3 orbitals are empty, this interaction is called donative with respect to the CO ligand. As in a Lewis base-Lewis acid interaction, a fraction of electronic charge is donated into the empty Ni orbitals from the doubly occupied CO 5r~ orbitals.

Weak covalent interactions with the d atomic orbitals are also present, leading to the electronic energy level scheme as shown in Fig. 4.4b. The total number of valence electrons is 18. The complex is stable at this electron count because of the large energy gap between highest occupied d levels and the antibonding (~

orbitals. This corresponds to the so-called 18 electron rule.

Compounds are stable as long as no significantly destabilized antibonding orbitals become occupied.

The dxy, d,x and d~y atomic orbitals on the Ni atom become part of the weakly destabilized antibonding levels due to interaction with the 5(~ CO levels. Since all of the d orbitals are doubly occupied, the overall 5r~ CO 4d Ni interaction is repulsive.

So far we have not discussed the interaction of the Ni electrons with the empty CO 2~* orbitals. Because of symmetry constraints, these orbitals interact with the doubly occupied d/2_y2and d ~2 orbitals (Figs. 4.4b and 4.4c). Bonding and antibonding orbitals are formed as a linear combination of the respective Ni d and

120 4 ~ B O N D I N G A N D E L E M E N T A R Y STEPS IN CATALYSIS

2 ~ ~

Fig. 4.5. Molecular orbital scheme of the electronic structure of CO.

CO 2~* orbitals. The bonding orbitals become doubly occupied and have mainly a d character. They are downwards shifted with respect to the Ni d x2_y2 and d ,2 orbitals.

The antibonding orbitals remain empty and have mainly a 2~* character. This interaction system is called the backdonation part of the CO-metal interaction, because the bonding interaction results in electron transfer from the Ni 4d atomic orbitals to the CO 2~* orbitals.

Since the 2~* orbitals of CO are antibonding with respect to the CO bond, population of the 2~* orbitals weakens the CO bond. This can be observed using infrared spectroscopy, where one usually finds a lowering of the CO stretching frequency when CO becomes part of co-ordination complexes or adsorbs to surfaces. It implies significant backdonation.

From the above it may be clear that the magnitude of this shift has no immediate relation to the strength of the metal to carbon bond. For instance, in Ni(CO)4 the bondstrength is dominated by the donative o-type interaction between the CO 5~ orbitals and Ni 4s and 4p atomic orbitals.

The orbital interaction scheme that describes bonding of ligands as CO, but also of other ligands, in terms of the sum of donative and backdonative interactions is called the Chatt-Dewar-Duncanson picture of the chemical bond.

Historically it was first used to describe bonding of ethylene. As we will see later, bonding to surfaces (Section 4.4.1.2) is quite analogous and there it is called the Blyholder model.

4 - - B O N D I N G A N D E L E M E N T A R Y STEPS I N C A T A L Y S I S 121 Within the classical Chatt-Dewar-Duncanson and Blyholder descriptions of chemical bonding as a sum of donative and backdonative interactions, only the attractive interactions (see Eqn. (4.12a)) between doubly occupied and empty valence orbitals are considered.

The case of Ni(CO)4 illustrates that repulsive interactions between doubly occupied orbitals, due to filling of antibonding orbitals are also important. The need to reduce this repulsive interaction often controls the particular bonding configuration of adatoms to surfaces or in cluster compounds.

For Co, having one electron less than Ni, one finds that C0(CO)4 can accept one additional electron. This is the reason why dimerization occurs giving the stable [C0(CO)412 complex, but bonding of C0(CO)4 with an hydrogen atom can also occur. Because of the stability of the C0(CO)4 anion (18 electron rule), the hydrogen atom in the hydrogen cobalt carbonyl complex has an acidic character:

HCo(CO)4 ~ H § + Co(CO)4- (4.16)

In contrast, a covalent metal hydride bond is formed in HMn(CO)s. Its bonding scheme is shown in Fig. 4.6. Manganese has seven valence electrons: the H atom donates one electron. Again, in the octahedral complex bonding with CO and H is dominated by the donative interaction between the doubly occupied CO 5c~

orbitals, the H s orbital and empty Mn 4s, 4p anddx2_y2 and the metald ,2 orbitals.

Similar to the situation as described in Section 4.2.1.2, bonding between H and Mn can be best described as electron transfer from Mn to H, to give Mn +~ and H- and a donative interaction between H - a n d empty Mn +1 orbitals. For simplicity in Fig. 4.6 the H(ls) orbital energy and Co 5c~ orbitals have been treated as equi- valent. One notes that HMn(CO)5 again is an 18 electron system. Now the antibonding

dx2_~2

and d~2 type orbitals are also non-occupied and the interaction between the d- as well as s, p type orbitals of Mn § and CO and H- is only of donative and attractive nature. It results in a stable complex, having stronger bonds than in Ni(CO)4.

Electron backdonation occurs between the doubly occupied d~y, dy~ and

d,x

orbitals and empty CO 2~* orbitals.

As a prelude to oxide or sulfide surfaces, it is useful to consider the electronic structure of Mn(CO)5. The removal of H changes octahedral co-ordination to one of fewer neighbours.

The resulting orbital scheme is shown in Fig. 4.7. Now only five occupied, bonding c~ type orbitals are formed and electrons redistribute over the non- and anti-bonding d-character orbitals of Mn. The Mn atom is neutral. The non-shared electron occupies an antibonding d ~2 type orbital (Fig. 4.7b) that has hybridized along the z-axis (the direction of the empty site), with s and p~ atomic orbitals. A dangling outwards directed orbital is formed containing one electron. Another dz 2, s,

p,

orbital combination is directed inwards and forms a strong bond with a CO 5c~ orbital.

122

i{

E /

3 /

/ /

/ ,...""'"

__ }~ ...

+ ~ +q-q- ::~i!ii

_{_}

..._

4 - - B O N D I N G A N D E L E M E N T A R Y S T E P S I N C A T A L Y S I S

. . . }_

...

\ \

dx2-~ dz2

. . . ~il~... "\

, "...

i ~',.. ~ \

4-++*

d~ dy, dz~ ' ^' ^'.. "... ^",."

9 .. ',

,+ } oco

++§

f r e e M n c o m p l e x f r e e C O

Fig. 4.6. (a) Structure of the HMn(CO)5 complex. (b) The molecular orbital scheme of octahedral HMn(CO)s complex; CO 2~* interaction not included.

We have seen that complexes become stabilized for particular electron counts.

When an increase in electron count results in the population of strongly anti- bonding orbitals, no stable compounds are formed. An analogous rule has been

4 - - B O N D I N G A N D E L E M E N T A R Y S T E P S I N C A T A L Y S I S

(a)

+ + + - 4 - -

/ ...

. ...

i..

, . . . . " " "

{ _7_ 0x, ~} ^,

... dz2 - - \

..,."

.. ... ~.

'... ~

~,z+ +~o} "..

~z.+ \\

^{9 .... -}....-

~:~' + 4 - + ++~o

...."

4- JJ

free atom complex free CO

123

(b)

CO

(c)

Fig. 4.7. (a) Molecular orbital scheme of Mn (C0)5. (b) The antibonding dz 2 type dangling bond orbital.

(c) The bonding d,2--5a orbital system.

124 4 w BONDING AND ELEMENTARY STEPS IN CATALYSIS

found for the stability and hence reactivity of metal clusters. It stems from the large overlap of the s and p atomic orbitals and the resulting importance of the s and p valence electrons to the cohesive energy of the metal clusters. Each atom can be considered to contribute one electron per atom to the s, p valence electron bonds of the cluster. In a closed packing geometry the clusters are approximately spherical, and hence, a stable cluster will be one with n = 2, 8, 18, etc., n being the number of atoms that are part of the cluster. Indeed, beam experiments indicate exceptional stability or nonreactivity for such cluster atom numbers for many metals. The shell rule, however, is very approximate and in most cases it is better to relate surface atom co-ordination numbers with cluster reactivity.

The description of the bonding of transition metal cations in oxide or sulfide surfaces, is very similar to that of the carbonyl complexes. The chemistry of these ions is related to that in co-ordination complexes when cation charge and local co-ordination are similar. We will illustrate this point of view for NiO and MoS2.

The local co-ordination of the cations is illustrated in Fig. 4.8.

In oxides or sulfides the anionic electrons are considered to contribute one doubly occupied o-type atomic orbital to the anion-cation bond and two doubly occupied ~-type orbitals.

The approach we follow is justified, as long as the anionic orbitals remain doubly occupied. Then the electron distribution around the anions is spheric.

Using this procedure the orbital schemes for bonding of octahedral Ni 2§ or prismatic M o 4+ c a n be readily deduced. The corresponding orbital schemes are shown in Fig. 4.9.

The Molecular Orbital scheme for NiO is very similar to that of HMn(CO)s.

The six bonding c~ type orbitals are mainly localized on oxygen or sulfur and doubly occupied. In NiO, as in HM(CO)4, the

dx2 y2

and d ~2 atomic orbitals are weakly antibonding. In MoS2 the splitting of the antibonding and non-bonding d-orbitals is different. Therefore at high d-electron-occupation the prismatic structure becomes less stable than octahedral co-ordination.

The stability of the compound increases when, compared to Ni 2§ the number of electrons on the central M ~+ ion decreases because less antibonding d-orbitals become occupied. Co B+ or Fe 2§ should give optimum bonding, since with less

Fig. 4.8. Local c a t i o n c o o r d i n a t i o n in N i O a n d MoS2.

4 - - BONDING AND ELEMENTARY STEPS IN CATALYSIS 125

~*{

A~

dx2.y2 dz2

q- +

d,,y dyz dzx

Ac

dzx dyz

I

dx2.y2 , . dxy

dz2

"{o ,++§ "{ +* }o**

+ +

+

a b

Fig. 4.9. (a) Molecular Orbital s c h e m e of o c t a h e d r a l l y c o o r d i n a t e d Ni 2+ in NiO. (b) M o l e c u l a r Orbital Scheme of prismatically c o o r d i n a t e d Mo 4§ in MoS2.

electrons the non-bonding or weakly bonding d-electrons become non occupied.

At the surface the cations and anions have fewer neighbours. The resulting electronic reorganization of valence electrons is quite analogous to that previous- ly discussed for Mn(CO)s.

In the (100) surface of NiO, Ni 2+ becomes five co-ordinated to oxygen. The outward-directed orbital is the rehybridized d~2 orbital, which is doubly occupied and hence will have a strong repulsive interaction with occupied adsorbate valence orbitals of similar symmetry.

Again Co B+ or Fe 2§ in similar co-ordination will have a strong interaction, because of the absence of this repulsive interaction.

Four co-ordinated Ni 2§ on edges of NiO will become more reactive because now also the occupied outward-directed

dyz

orbitals can backdonate electrons into empty asymmetric adsorbate orbitals. As we will see later, this interaction is responsible for dissociative adsorption reactions. The corresponding orbital scheme and structure is shown in Fig. 4.10.

This very short intermezzo on the electronic structure of transition metal compounds had as its main purpose to illustrate that bonding is very similar to that in molecular co-ordination compounds. Hence their chemical reactivity can be understood using the same concepts.

The presence of Lewis basic oxygen or sulfur anions around the positively charged cation provides additional bonding positions for protons generated in heterolytic dissociation reactions (Fig. 4.11). The cation acts in essence as a Lewis

126

0 Ni 0 x

/

0

4 - - BONDING AND ELEMENTARY STEPS IN CATALYSIS

(3'

a * {

d,~

-dx'.~

+

^d~dz2 ^z&a

4-

':{ ++++ }

Fig. 4.10. (a) Local s t r u c t u r e of 4 c o o r d i n a t e d N i 2§ in N i O surface. (b) O r b i t a l s c h e m e of N i 2§

c o r r e s p o n d i n g to s t r u c t u r e 4.10a.

Co ~ 0

.0

H~ 6+

, ~ H

J 9

i Oo

' 9

i i i

/ co2§ - O - c o

9 Co

I

5 - 6+

H H

I I

0 ,.Co O - - C o 0

I

Fig. 4.11. H e t e r o l y t i c d i s s o c i a t i o n of H2 o n a Co 2+ centre.

acid. The hydrogen binds to it as an anion, as in HM~(CO)5. Its acidity depends strongly on the electrostatic field generated by its neighbour anions.

To illustrate oxidative addition of H2 to a metal centre and the importance of the rehybridization due to the interaction with metal atom ligands, we will compare the reaction of H2 by Ni atoms and the Ni(PH3)2 complex [3]. Oxidative addition is the analogue of dissociative adsorption of a molecule to a catalyst centre. A terminology used in surface science and heterogeneous catalysis.

In the Ni atom, the 4s orbital contains two electrons. This spatially extended, doubly occupied orbital experiences a large repulsive interaction with the doubly

4 B BONDING AND ELEMENTARY STEPS IN CATALYSIS 127

occupied r~ orbital of H2 of the same symmetry. Dissociation of H 2 has a large activation energy because electron promotion has to occur from 4s to 3d atomic orbitals in order to obtain dissociation. Co-ordination of ligands, as in Ni(PH3)2, decreases the promotion energy. The Ni(PH3)2 complex has initially a linear geometry (Fig. 4.12).

Directed occupied bonding orbitals and non-occupied antibonding orbitals between the Ni atom and the two phosphorus ligands are formed when two electrons of the Ni 4s orbitals are promoted into the Ni 3d orbitals. The Ni 4s and 4p atomic orbitals hybridize into two empty atomic orbitals directed to the occupied lone pair on the phosphorus atom. When the hydrogen molecule approaches this complex, the interaction is again weak. The empty, spatially extended Nip orbitals required for bonding have been pushed upwards to a high

H3P-Ni-PH3

(a)

Fig. 4.12. The interaction between a Ni atom and a h y d r o g e n molecule. (a) Geometry of H3P-Ni-PH3.

(b) The orbital scheme of H3P-Ni-PH3.

128 4 -- BONDING AND ELEMENTARY STEPS IN CATALYSIS

PH3 (C) PH3

d; )

(d) Q

Fig. 4.12 continued. (c) Orbitals in the square-planar complex H2Ni(PH3)2. (d) Rehybridized orbitals in H2Ni(PH3)2.

energy. Dissociation occurs only when the Ni(PH3)2 molecule is deformed in such a way that empty spatially extended orbitals become available.

When the HBP-Ni-PH3 angle becomes 90 ~ hybridization between the dxz and s atomic orbitals, rather than s-p hybridization, results in directed orbitals suitable for co-ordination with the PH3 ligands. In the process, empty antibonding orbitals having an s and dx~ nature are created, as sketched in Fig. 4.12d. Their orientation and symmetry are suitable for interaction with the hydrogen molecule. Now the antibonding H2 o orbital can be populated and the molecule will dissociate. Note that hybridization on the Ni atom is n o w d9s 1, which is very similar to the hybridization of a Ni atom in the Ni metal surface.

For dissociative adsorption of H2 we pointed out the importance of reducing repulsive interactions between doubly occupied orbitals and the need for easily accessible occupied orbitals, asymmetric with respect to the reaction co-ordinate, that backdonate electrons into antibonding adsorbate orbitals. Occupation of the latter weakens the adsorbate bonds and reduces the activation energy for dissociation.

In organometallic chemistry dissociative adsorption is called oxidative addition. The hydrogen atoms formally become negatively charged. In the

4 - - B O N D I N G A N D E L E M E N T A R Y STEPS I N C A T A L Y S I S 129 Ni(PH3)2 complex, the valence of Ni changes from zero to 2+ in the H2Ni(PH3)2 complex.

Whereas in oxidative addition electrons are used to form new bonds between metal centre and adsorbate fragment, in the reverse reaction reductive elim- ination of the adsorbate donates electrons to the metal centre. The H-atoms recombine to form H2 that desorbs.

Recombination reactions in which bonds are formed between fragments around a metal centre, even when no desorption occurs, also require the avail- ability of empty transition metal orbitals, asymmetric with respect to the reaction co-ordinate.

We will illustrate this for the insertion reaction of CO into a CH3 group co-ordinated to Pd 2+. This reaction can be considered as a prototype for C-C bond formation catalyzed by a metal centre (see Fig. 4.13).

The formal charge on the methyl group is -1, which implies that 2 electrons occupy the o-symmetric lone pair orbital. When the CH 3 fragment migrates to CO and a carbon-carbon bond is formed, the doubly occupied 5r~ CO orbital interacts with the doubly occupied (~--CH3 orbitals resulting in doubly occupied bonding and antibonding orbitals, giving a repulsive interaction (Fig. 4.13).

The empty

dx2_y2

orbital on the Pd atom can accept electrons from the antibonding C-C fragment orbital, thus reducing the repulsive interaction [4,5].

As we will also see later, the possibility of stabilizing surface interactions with occupied or unoccupied orbitals, asymmetric with respect to the reaction co- ordinate, will also be essential to enable surface associative or dissociative reactions.

Dalam dokumen Catalysis: An Integrated Approach (Halaman 121-133)