Periodic Trends in the Bond Energies of Transition Metal Complexes

---

Chapter 19

Periodic Trends in the Bond Energies of Transition Metal Complexes Density Functional Theory

Downloaded by UNIV OF AUCKLAND on May 3, 2015 | http://pubs.acs.org Publication Date: June 25, 1990 | doi: 10.1021/bk-1990-0428.ch019

Tom Ziegler and Vincenzo Tschinke Department of Chemistry, University of Calgary, Calgary, T2N 1H4, Canada Accurate Density Functional calculations make possible today the systematic investigation of periodic trends in the bond energies of transition metal complexes. Computational results are presented for metal-metal bonds in dimers of the group 6 transition metals, metal-ligand bonds in early and late transition metal complexes, and metal-carbonyl bonds in hexa- penta- and tetra-carbonyl complexes.

The dearth of reliable experimental data on bond dissociation energies is felt throughout the field of organometallic chemistry. Accurate theoretical studies should afford a much needed supplement to the sparse available experimental data on metalligand bond energies, necessary for a rational approach to the synthesis of new transition metal complexes. Recently, Density Functional investigations of molecular bond energies have gained novel impetus due to the introduction by Becke (7) of a gradient correction to the Hartree-Fock-Slater local exchange expression, -1

IV, t^LSD/NL

C

HFS V f t

2

PÏ(?i)l

ï

i l+YB

[p][(ri)]

7/3

[p?(ri)]

δΓ!

(1),

8 / 3

where is a spin density and ββ and γβ are parameters. In conjunction with appropriate approximations for antiparallel spin correlations, the expression of Eq.(l) provides near-quantitative estimates (7) of bond energies in main-group compounds. In this contribution we shall present several applications of the new method, which we shall refer to as LSD/NL, to the calculation of bond energies of transition metal complexes. We shall focus on trends along a transition period and/or down a transition triad. The following subjects will be discussed: a) metal-metal bonds in dimers of the group 6 transition metals; b) metal-ligand bonds in early and late transition metal complexes; c) the relative strength of metal-hydrogen and metalmethyl bond in transition metal complexes; d) the metal-carbonyl bond in hexa- pentaand tetra-carbonyl complexes. 0097-6156/90A)428-0279$06.00/0 © 1990 American Chemical Society In Bonding Energetics in Organometallic Compounds; Marks, T.; ACS Symposium Series; American Chemical Society: Washington, DC, 1990.

280

BONDING ENERGETICS IN ORGANOMETALLIC COMPOUNDS

Table 1 Calculated bond energies [D(M-M)] (eV) and metal-metal bond distances (RM-M) (Â) for Cr , M02 and W2 2

D(M-M)

Cr

2

M02

w

RM-M 11

Calc. 1.75 4.03 4.41(3.54)

Exp. 1.5610.2 4.18±0.2

-

a

2

11

Exp. 1.69 1.93

Calc. 1.65 1.95 2.03(2.07)

-

a

Downloaded by UNIV OF AUCKLAND on May 3, 2015 | http://pubs.acs.org Publication Date: June 25, 1990 | doi: 10.1021/bk-1990-0428.ch019

Non-relativistic results.

Computational Details In the present set of calculations we have used the functional proposed by Becke (7), which adopts a non-local correction to the HFS exchange, and treats correlation between electrons of different spins at the local density functional level. All calculations presented here were based on the LCAO-HFS program system due to Baerends et al. (2) or its relativistic extension due to Snijders et al.(3), with minor modifications to allow for Becke's non-local exchange correction as well as the correlation between electrons of different spins in the formulation by Stoll et al. (4) based on Vosko's parametrization (5) from homogeneous electron gas data. Bond energies were evaluated by the Generalized Transition State method (6), or its relativistic extensions (7). A double ζ-STO basis (8) was employed for the ns and np shells of the main group elements augmented with a single 3d STO function, except for Hydrogen where a 2p STO was used as polarization. The ns, np, nd, (n + l)s and (n + l)p shells of the transition metals were represented by a triple ζ-STO basis (3). Electrons in shells of lower energy were considered as core and treated according to the procedure due to Baerends et al. (2). The total molecular electron density was fitted in each SCFiteration by an auxiliary basis (9) of s, p, d, f and g STOs, centred on the different atoms, in order to represent the Coulomb and exchange potentials accurately.

Metal-Metal Bond Strength of the Dimers Cr , M02 and W 2

(10)

2

The bond energies (77) of the metal dimers C r and Mo are accurately known experimentally and we note that several theoretical accounts of the bonding in these systems have already appeared, based on ab initio (11a) and Density Functional Theory (72). However, no calculation has been reported for W , nor are there any experimental data available. The bond energies in Table 1 were calculated by evaluating the energy difference 2

2

2

ΔΕ = 2EÇS) - E(M )

(2)

2

7

l

between two metal atoms in the S state corresponding to the nd$(n + l)s configuration, and M . For Cr and Mo , ΔΕ represents the bond energies D(Cr-Cr) and D(Mo-Mo), respectively, since Cr and Mo have a spherical 'S ground-state. However, the W-atom has a D ground-state with the configuration 5d 6s , thus we have subtracted for W the experimental energy difference (.37 eV) (75) between the D and the S states twice to arrive at D(W-W) of Table 1. 2

2

2

5

4

2

2

5

7

In Bonding Energetics in Organometallic Compounds; Marks, T.; ACS Symposium Series; American Chemical Society: Washington, DC, 1990.

19. ZIEGLER & TSCHINKE

Table 2

Periodic Trends in the Bond Energies

281

1

Calculated [D(M-L)] bond energies (kJ mol") and optimized (RM-L) bond distances

(A) in CI3ML

L H CH

Ti 250.7 267.5 210.9 453.2 426.9 293.3 364.7 190.6 410.4

3

S1H3

OH Downloaded by UNIV OF AUCKLAND on May 3, 2015 | http://pubs.acs.org Publication Date: June 25, 1990 | doi: 10.1021/bk-1990-0428.ch019

OCH3

SH NH PH CN

D(M-L) M Zr 297.2 309.5 239.5 527.2 484.5 347.9 420.6 225.6 457.6

2

2

a

Hf 313.5 326.6 272.2 535.9 506.6 360.1 439.1 233.9 477.9

Ti 1.70 2.13 2.63 1.83 1.86 2.28 1.87 2.24 2.06

RM-L M Zr 1.82 2.26 2.78 1.95 1.99 2.47 2.01 2.48 2.21

Hf 1.80 2.25 2.79 1.97 2.01 2.47 2.04 2.47 2.23

a

Optimized from a quadratic fit through three energy points corresponding to three different M-L distances. The calculated bond energies and equilibrium bond distances RM-MforCr and M02 are in good accord with experimental values, as can be seenfromTable 1. In contrast to other calculations based on DFT, we have employed in the present work (n + l)f polarization functions. Their contribution to the bond energies are modest, 0.2 - 0.4 eV. On the other hand, the contributions to D(M-M) from the non-local correction to the exchange are -1.8 and -2.4 eV for M02 and Cr , respectively, and are thus important in determining the agreement with experiment. We predict that W , after the inclusion of relativistic effects, should have a stronger metal-metal bond than M02. Even in the non-relativistic case, the bonding interaction is stronger in W than in Mo if the two metal atoms are referred to the same S reference state. 2

2

2

2

2

7

Metal-ligand Bond Strengths in the Early Transition Metal Systems CI3ML and Late Transition Metal Systems LCo(CO) (14) 4

The way in which metal-ligand bond energies of early transition metals and f-block elements differ from those of middle to late transition metals, or metal-ligand bond energies of 3d and 4f elements differ from those of their heavier congeners, has been the subject of many experimental (75) as well as a few theoretical studies (76) over the past decade. We shall present here calculations on the D(M-L) bond strength in the CI3ML (1) model systems of the early transition metals M = Ti, Zr and Hf, as well as the LM(CO)4 model system with the late transition metal M = Co, for a number of rudimentary ligands, L = H, C H , S1H3, OH, SH, O C H 3 , N H , P H and CN. The calculated bond energies D(M-L) are displayed in Table 2. The ligands L = OH, O C H 3 , with the coordinating atoms of the highest electronegativity and the most polar CI3M-L bond have the largest D(M-L) bond energies. The ligand L = S1H3, with the coordinating atom of the lowest electronegativity and the least polar M-L bond, has a modest D(M-L) bond energy. For the series of ligands NH , SH, C H 3 3

2

2

2

In Bonding Energetics in Organometallic Compounds; Marks, T.; ACS Symposium Series; American Chemical Society: Washington, DC, 1990.

BONDING ENERGETICS IN ORGANOMETALLIC COMPOUNDS

282

f

Ev 3

2e

θ

ib —

i£*

2

1a,

^ l a 1

θ

1

θ

SH > CH3 > H. 2

CI

\

.M I,

CI CI It is clear from Table 2 that zirconium, and even more so hafnium, form stronger M-L bonds to the ligands under investigation than titanium in the CI3M-L systems. We have found that the calculated increase in D(M-L) down the triad is primarily caused by a corresponding increased overlap between the singly occupied lai-orbital of CI3M and the singly occupied orbital on the ligand L, see Figure 1, which is in turn responsible for an increased σ-bonding interaction between the metal centre and the ligand (77). There are few thermochemical data available for M-L bond involving group 4 metals. To our knowledge, of the ligands under consideration data (15a) are only available for L = CR , NR and OR (R = alkyl), for the homoleptic M(CR )4, M(NR )4, and M(OR)4 systems with M = Ti, Zr and Hf. The M-L bonds in these systems follow the same trend as observed here for the CI3M-L systems, with the bond energy increasing down the triad as well as with the increasing electronegativity of the corresponding atom on the ligand (Ο > Ν > C). Group 4 metals are known (18) to form several complexes involving M-L bonds with L = S1R3, PR , and SR ligands. However, the corresponding D(M-L) bond energies have not been determined experimentally. Perhaps not surprisingly, we find that the M-L bonds of L = S1H3, P H , and SH are weaker and less polar than the M-L bonds of the homologous ligands L = C H , N H , and OH (Table 2). Comparative experimental data on M-H and M-CH3 bond energies of early transition metals are not available. However, it has been asserted (79) that the M-H and M-CH3 bond strengths of early transition metal complexes are quite similar. Indeed, we find D(M-H) and D(M-CH ) to be quite similar in the CI3M-H and CI3MCH3 systems, respectively. We shall dedicate the next section to the relative strengths 3

2

3

2

2

2

3

2

3

In Bonding Energetics in Organometallic Compounds; Marks, T.; ACS Symposium Series; American Chemical Society: Washington, DC, 1990.

19. ZIEGLER & TSCHINKE

Periodic Trends in the Bond Energies

283

Table 3 Calculated [D(Co-L)] bond energies (kJ mol") and optimized (RCo-L) bond distances (Â) in LCo(CO)4 1

L H CH SiH OH

Downloaded by UNIV OF AUCKLAND on May 3, 2015 | http://pubs.acs.org Publication Date: June 25, 1990 | doi: 10.1021/bk-1990-0428.ch019

3

3

DCC6-L) 230. 160. 211.6 232.4

RCO-T.

1.55 2.11 2.73 2.09

L SH NH PH CN

D(Co-L) 168.9 145.5 145.5 304.3

2

2

RCo-I,

2.49 2.09 2.43 2.04

of the M-H and M-CH3 bonds in systems ranging from early transition metal and fblock element complexes to middle and late transition metal complexes.

CO OC CO

We shall now consider the M-L bond energies in the late transition metal complexes LCo(CO)4 (2). The σ-bond in LCo(CO)4 is considerably less polar than in CI3ML. This is primarily so because cobalt is more electronegative than the group 4 metals Ti, Zr and Hf, and as a consequence the lai metal based frontier orbital, involved in the σ-bond, is of lower energy than the frontier orbital lai of CI3M (see Figure 1). For ligands other than H we had in the CI3ML systems favourable donor-acceptor interactions from occupied ligand orbitals to the two empty e-sets, see Figure 1. In Co(CO)4 the metal based d-orbitals of e-symmetry are fully occupied and the corresponding interactions between occupied ligand orbitals and either le or 2e are as a consequence repulsive. As a consequence, we find in accord with available experimental evidence (20) that all ligand except hydrogen form stronger bonds to the early transition metal Ti than to the late transition metal Co, see Table 2 and 3.

The Relative Strengths of the Metal-Hydrogen and the MetalMethyl Bonds in Transition Metal Complexes (14,21,22). The breaking or formation of metal-hydrogen and metal-alkyl bonds is an integral part of most elementary reaction steps in organometallic chemistry. As a consequence, considerable efforts have been directed toward the determination of M-H (15b) and Malkyl bond strength (23) as a prerequisite for a full characterization of the reaction enthalpies of elementary steps in organometallic chemistry.

In Bonding Energetics in Organometallic Compounds; Marks, T.; ACS Symposium Series; American Chemical Society: Washington, DC, 1990.

BONDING ENERGETICS IN ORGANOMETALLIC COMPOUNDS

Downloaded by UNIV OF AUCKLAND on May 3, 2015 | http://pubs.acs.org Publication Date: June 25, 1990 | doi: 10.1021/bk-1990-0428.ch019

284

Fig. 2 Calculated M-H and M-CH3 bond energies for several actinide and transition metal complexes.

As already mentioned, the strengths of M-H and M-Alkyl bonds are comparable for early transition metals (19). According to sparse experimental data, the same trend is observed in actinide complexes (24). By contrast, data for alkyl (20a,15e,25,26) and hydride (20e,26,27) complexes of middle to late transition metals indicate that the M-H bond is stronger than the M-Alkyl bond by some 40-80 kJ mol" . This difference in strength has implications for the relative ease by which ligands can insert into the M-H and M-Alkyl bonds (28). Also, it is one of the thermodynamic factors, along with the relative order of the bond energies H < H Alky < Alkyl-Alkyl, which favour the oxidative addition (19,29) to metal centres of H compared to H-Alkyl and Alkyl-Alkyl bonds. In the preceding section we have presented results on the relative bond strengths of the M-H and M-CH3 bonds of model complexes CI3M-R (R = H, CH3) involving the early transition metals Ti, Zr and Hf, as well as the late transition metal complex R-Co(CO)4. Here, we shall present additional results on actinide metal complexes as well as on middle and late transition metal complexes, in an attempt to supplement the rather sparse experimental data available. We have conducted calculations on the bond energy D(M-R) (R = H, CH3) of the model actinide complexes CI3M-R (M = Th and U) as well as the middle transition metal complexes R-M(CO)5 (M = Mn, Tc and Re) and the late transition metal complexes R-M(CO)4 (M = Co, Rh and Ir). These results are depicted in Figure 2, along with the corresponding results presented in the preceeding section. The results depicted in Figure 2 seem to indicate that the trend which assigns comparable M-H and M-CH3 bond strengths in early transition metal and actinide complexes but a stronger M-H bond in middle and late transition metal complexes, is 1

2

2

In Bonding Energetics in Organometallic Compounds; Marks, T.; ACS Symposium Series; American Chemical Society: Washington, DC, 1990.

19. ZIEGLER & TSCHINKE

Periodic Trends in the Bond Energies

285

of general validity. It also appear that both the M-H and the M-CH3 bond strengths increase down a triad for transition metal complexes.

Downloaded by UNIV OF AUCKLAND on May 3, 2015 | http://pubs.acs.org Publication Date: June 25, 1990 | doi: 10.1021/bk-1990-0428.ch019

2a

2h

The reduced strength of the M-CH3 bonds in middle and late transition metals can be readily explained in terms of destabilizing three- and four-electron two-orbital interactions which occur between the fully occupied 1σ (3a) and 1π (3b) methyl orbitals and the fully or singly occupied d-orbitals of matching symmetries present on the metal centres. By contrast, in early transition metal and actinide complexes, the metal orbitals of π-symmetry are vacant and are involved in stabilizing interactions with 2k A destabilizing three-electron two-orbital interaction between 2â and the single-occupied metal orbital lai (Figure 1) is still present. However, the M-CH3 bond is more polar in early transition metal complexes than in middle and late transition metal complexes and as a consequence electronic charge is transferred from the metal orbital lai to the methyl ligand, thereby relieving the destabilizing interaction between lai itself and 3a. The comparable strengths of the M-H bonds in the complexes studied is perhaps not too surprising, since H is a simple one-orbital ligand without additional occupied orbitals involved in four-electron two-orbital interactions or π-donor-acceptor interactions. Finally, the increase in strength of both the M-H and M-CH3 bonds down a triad, see Figure 2, is primarily related to an increase in the overlap between the metal lai orbitals and the matching ligand orbitals, which leads to a stronger σ interaction. Such an increase in overlap occurs as the metal d-orbitals become more diffuse down the triad. A comparison between our calculated results for D(M-R) (R = H, CH3) and the few available experimental data is presented in Table 4. We find in general a good agreement with the experimental bond energies. Also, the stability order D(M-L) > D(M-CH3) in middle and late transition metal complexes supported by our theoretical study is consistent with data on organometallic reactions in which M-L and M-CH3 bonds are formed or broken. Thus, CO will readily insert into a M-CH3 bond whereas the corresponding insertions into M-H bonds are virtually unknown (28), and methyl has likewise a larger migratory aptitude toward most other ligands than hydride. The H2 molecule is known to add oxidatively and exothermically to several metal fragments where the corresponding oxidative additions of the H-Alkyl and Alkyl-Alkyl bonds are unknown and probably endothermic as a consequence of the weak M-R bond (19). At this stage, it is important to comment on the role of the non-local correction to the exchange in the calculated bond energies. For middle and late transition metal complexes, the non-local correction reduces significantly the D(M-CH3) values (by 105 kJ mol" for CH3-Mn(CO)5) whereas the corresponding D(M-H) bond energies are decreased to a lesser extent (by 13 kJ moF for H-Mn(CO)s). Thus, it is apparent that Becke's non-local correction to the exchange is essential to assure the good agreement of the LSD/NL results with experiment, whereas the FTFS and the LSD methods not only tend to give too large bond energies, but in some cases predict the wrong order for the M-H and M-CH3 bond strengths. 1

In Bonding Energetics in Organometallic Compounds; Marks, T.; ACS Symposium Series; American Chemical Society: Washington, DC, 1990.

BONDING ENERGETICS IN ORGANOMETALLIC COMPOUNDS

286

Table 4 (kJ mol')

Calculated and experimental values for the bond energies D(M-R) (R = H, CH3)

1

M-H Cl Th-H 3

CI3U-H CI3T1-H

Cl Zr-H Cl Hf-H H-Mn(CO) H-Tc(CO) H-Re(CO)5 H-Co(CO)4 H-Rh(CO) H-Ir(CO)

Downloaded by UNIV OF AUCKLAND on May 3, 2015 | http://pubs.acs.org Publication Date: June 25, 1990 | doi: 10.1021/bk-1990-0428.ch019

3

3

5

5

4

4

a c

LSD/NL 318.0 293.3 250.7 297.2 313.5 225 252 282 230 255 286

LSD/NL 333.9 CI3U-CH3 302.1 CI3T1-CH3 267.5 Cl Zr-CH 309.5 Cl Hf-CH 326.6 CH-Mn(CO)5 153 178 CH-Tc(CO)5 200 CH-Rc(CO)5 160 CH -Co(CO) 190 CH -Rh(CO) 212 CH -Ir(CO)

Exp. ~335. 319.7

Exp. ~335. 302.9 — —

M-CH3

a

Cl Th-CH

a

3

a

— b

213 — — 238 — —

3

3

3

3

3

a

b

153 — — — — —

3

3

3

c

3

4

3

4

3

4

Experimental bond energies from Ref. 24 correspond to Cp2 MC1-R systems. Ref.27e.

b

Ref. 20c.

Thermal Stability and Kinetic Lability of the Metal-Carbonyl Bond (30) The extensive use of coordinatively saturated mono-nuclear carbonyls as starting materials in organometallic chemistry, along with their volatility and high molecular symmetry, has prompted numerous experimental (15a,31,32,33) and theoretical (34,35) studies on their structure and reactivity. Special attention has been given to the degree of σ-donation and π-back-donation (34b-g,35a,35e) in the synergic (34k) M-CO bond. However, in spite of many experimental (32) investigations, there is still a lack of basic data on the thermal stability and kinetic lability of the M-CO bond in essential metal carbonyls such as M(CO) (M = Cr, Mo, W), M(CO) (M = Fe, Ru, Os) and M(CO)4 (M = Ni, Pd, Pt), particularly with respect to the carbonyls of the secondand third-row metals. Theoretical methods have begun to play a role in determining the energetics of organometallics (35g) and ab initio type methods have recently been applied to calculation on the M-CO bond strength of Cr(CO) (35d-e), Fe(CO) (35a-cf), and Ni(CO)4 (35a J), but not yet to M-CO bond strength of their second- and third-row homologues. Here, we shall present LSD/NL calculations on the intrinsic mean bond energy D(M-CO) and first CO dissociation energy ΔΗ of Cr(CO)e, Fe(CO)s, and Ni(CO)4 as well as their second- and third-row homologues. We shall here be concerned with periodic trends in the strength of the M-CO bonding interaction within the triads M = Cr, Mo, W; M = Fe, Ru, Os; and M = Ni, Pd, Pt. As measures for the M-CO bonding interaction in the hexacarbonyls (4a). pentacarbonyls (4b) and tetracarbonyls (&), we will consider the intrinsic mean bond energy D(M-CO) between M (in its low-spin valence state) and the η CO ligands, as well as the bond energy ΔΗ between M(CO) -i and CO. 6

5

6

5

n

In Bonding Energetics in Organometallic Compounds; Marks, T.; ACS Symposium Series; American Chemical Society: Washington, DC, 1990.

19. ZIEGLER & TSCHINKE

Periodic Trends in the Bond Energies

287

Table 5 Comparison between calculated and experimental values for the mean bond energy Ε of several carbonvl complexes. Calculated values for D(M-CO) and AE are also given. All values in kJ mol" prep

1

M(CO) E(Exp.) Cr(CO) Mo(CO) W(CO) Fe(CO) Ni(CO)

D(M-CO)

n

c

1/π·ΔΕρ ρ

E

100.7 51.6 54.4 98.42 —

110 126 156 117 —

Γβ

c

211 178 210 216.8 178.9

6

6

6

5

Downloaded by UNIV OF AUCKLAND on May 3, 2015 | http://pubs.acs.org Publication Date: June 25, 1990 | doi: 10.1021/bk-1990-0428.ch019

4

a

b

Réf. 15a. Experimental intrinsic mean dissociation energy D(M-CO), Réf. 15a. dissociation energy. Ο

a

110 151 179 118.4 191 a

a

a

b

c

Mean bond

Ο

There are two sets of experimental data with a bearing on the M-CO bond strength in M(CO) , namely, the mean bond energy Ε corresponding to the process n

M(CO) (g) -> M(g) + AzCO(g) - nE n

(2a)

and the first bond dissociation energy ΔΗ corresponding to the process M(CO) -> M(CO) -i + CO - ΔΗ n

n

(2b).

It is important to note that Ε is given by E = D(M-CO)-l/ME

prep

(3),

where A E is the energy required to promote the metal atom from its high-spin electronic ground state to the low-spin valence configuration. As a consequence, one can not conclude that the order of Ε will correspond to the order of D(M-CO) down a triad, since ΔΕρ™ might differ significantly for the three elements. The first bond dissociation energy ΔΗ is on the other hand a direct measure for the strength of the M-CO bond interaction. It is further an extremely important kinetic parameter, since the dissociation process of Eq.(2b) is assumed to be a key step in the large volume of kinetically useful substitution reactions (36) p r e p

M(CO) + L -» M(CO) -1L + CO n

n

(4)

where L is introduced into the coordination sphere of M by replacing one carbonyl ligand. Our computational results for D(M-CO) and ΔΗ are depicted in Figure 3. It appears from the values of D(M-CO) (Figure 3a) that strength of the M-CO bond

In Bonding Energetics in Organometallic Compounds; Marks, T.; ACS Symposium Series; American Chemical Society: Washington, DC, 1990.

288

BONDING ENERGETICS IN ORGANOMETALLIC COMPOUNDS

D(M-CO) kJ mol

Δ

•1

Η

kJ mol"

Downloaded by UNIV OF AUCKLAND on May 3, 2015 | http://pubs.acs.org Publication Date: June 25, 1990 | doi: 10.1021/bk-1990-0428.ch019

A 250

-

250

200

-

200

150

-

150

-

100

-

100

-

50

-

50

-

M(CO)

M(CO)

4

5

M(CO)

6

H

M(CO)

4

M(CO)

5

M(CO)

6

Fig. 3 Calculated bond energies for M(CO)6 (M = Cr, Mo, W), M(CO)5 (M = Fe, Ru, Os), and M(CO)4 (M = Ni, Pd, Pt): a) intrinsic bond energies D(M-CO); b) first CO dissociation energies ΔΗ. decreases going from the 3d to the 4d metal of a triad, to increase again for the complex of the 5d metal. The destabilization of the M-CO bond of the 4d and 5d elements, compared to their 3d counterparts increases in going from the Cr to the Ni group; in fact within the group 6 carbonyls, the M-CO bond in the W carbonyl is at a par in strength with the bond in the Cr complex. Among the systems studied, the strongest M-CO bond is assigned to the Fe(CO)5 complex, whereas the weakest bond is found in the Pd(CO)4 complex. It is important to note that the first dissociation bond energies ΔΗ (Figure 3b) follow closely the same trends observed for the intrinsic bond energies D(M-CO) (Figure 3a).

ο v

co

C

Ο

σ ill

The periodic trends discussed above can be readily rationalized in terms of the stabilizing electronic interactions and the destabilizing steric interactions which determine the strength of the M-CO bond. The electronic terms are represented by πback-donation from occupied nd-orbitals on the metal centre to the empty n orbital (5a). as well as σ-donation from the doubly-occupied Oco orbital (£10 to vacant nd orbitals. The steric terms are dominated by the repulsive four-electron two-orbital interactions between σ and occupied nd orbitals on the metal centre. Our results show that electronic factors are most favourable for the pentacarbonyls where both πc o

In Bonding Energetics in Organometallic Compounds; Marks, T.; ACS Symposium Series; American Chemical Society: Washington, DC, 1990.

19. ZIEGLER & TSCHINKE

289

Periodic Trends in the Bond Energies

Table 6 Comparison between calculated and experimental values for the first CO dissociation energy ΔΗ, values in kJ mol". Calculated values do not include geometry relaxation of the fragments M(CO) _i 1

n

M(CO) Cr(CO) Mo(CO) W(CO) n

6

6

6

a

LSD/NL 147 119 142

Exp. 162 126 166

M(CO) Fe(CO) Ru(CO) Ni(CO) n

a

a

b

5

a

4

b

LSD/NL 185 92 106 e

b

5

Exp. 176 117 104 d

e

f

c

Ref. 37. Equatorial CO dissociation energy. The dissociation product is the fragment Fe(CO)4 in its singlet state. Réf. 32a. Ref. 38. Ref. 39. Downloaded by UNIV OF AUCKLAND on May 3, 2015 | http://pubs.acs.org Publication Date: June 25, 1990 | doi: 10.1021/bk-1990-0428.ch019

d

e

f

back-donation and σ-donation are important, whereas the steric interactions are most favourable for the M-CO bond among the hexacarbonyls, where all nd orbitals of σsymmetry are empty and only mild repulsive interactions between Geo and the occupied ns and np metal orbitals are present. For first-row transition metals, the repulsive interactions between occupied nd- and Gco-orbitals are still modest, since the nd-Gco overlap integrals are relatively small for the contracted 3d-orbitals compared to the more diffuse 4d and 5d orbitals. Thus, electronic factors will make the intrinsic mean bond energy larger for Fe(CO)s than for Cr(CO>6 (Figure 3a). On the other hand, in carbonyls of 4d and 5d elements, where repulsive interactions between occupied nd- and Cco-orbitals are considerable, the steric factors cause the M-CO bonds in Ru(CO)5 and Os(CO)s to be weaker than in Mo(CO)6 and W(CO)6, respectively. The tetracarbonyls, in which all interactions between the nd- and acoorbitals are repulsive, have weaker M-CO bonds than the corresponding hexacarbonyls and pentacarbonyls in each of the transition series (Figure 3a). Finally, relativistic effects will stabilize the 5d element carbonyl compared to the 4