valence bond description of the co ligand

• In the CO molecule both the C and the O atoms are sp hybridized.

• The singly occupied sp and pz orbitals on each atom form a σ and a π bond,

respectively.

• This leaves the C py orbital empty, and the O py orbital doubly occupied, and so the

second π bond is formed only after we have formed a dative bond by transfer of the

lone pair of O py electrons into the empty C py orbital.

Valence Bond description of the CO ligand

• This transfer leads to a Cδ−−−−−Oδ+• This transfer leads to a Cδ−−−−−Oδ+

polarization of the molecule, which is

almost exactly canceled out by a partial

Cδ+−−−−Oδ− polarization of all three

bonding orbitals because of the higher

electronegativity of oxygen.

• The free CO molecule therefore has a

net dipole moment very close to zero.

C Osp sp spsp

p

p p

pdative

1

MO correlation diagram for CO

2p

C

HOMO

LUMO

OC

2p

O

The CO LUMO orbitals are anti bonding (π*). These

are empty orbitals and accept electron density

from the metal centre via π-backbonding with the

metal dxy, dxz and dyz orbitals.

2s

2s

2p

bond order = (8 - 2)/2 = 3

The CO HOMO orbital is a bonding orbital of σsymmetry with significant electron density on the

carbon. This orbital forms a σ bond with metal p , dz2

and dx2–y2 orbitals. This is a filled orbital and

donates electron density to the metal centre.

2

LUMO

∆∆∆∆

T2g3 x

dx2-y2dz2

HOMO

∆∆∆∆O

3

• The net result is that C becomes more positive on coordination, and O becomes more negative.

This translates into a polarization of the CO on binding.

• This metal-induced polarization chemically activates the CO ligand making C more sensitive to

nucleophilic attack and O more sensitive to electrophilic attack.

• The polarization will be modulated by the effect of the other ligands on the metal and by the

net charge on the complex.

� In LnM(CO), the CO carbon becomes particularly δ+ in character if the L groups are good π

acids or if the complex is cationic, e.g. Mo(CO)6 or [Mn(CO)6]+, because the CO-to-metal σ-

donor electron transfer will be enhanced at the expense of the metal to CO back donation.

� If the L groups are good donors or the complex is anionic, e.g. Cp2W(CO) or [W(CO)5]2−,

back donation will be encouraged, the CO carbon will lose its pronounced δ+ charge, but

the CO oxygen will become significantly δ−.

• The range can be represented in valence bond terms the extreme in which CO acts as a pure σdonor, through to the extreme in which both the π∗x and π∗y are both fully engaged in back

bonding.

5

CO's render the electron rich Cr metal electrophilic via strong π-backbonding. Complexation of benzene with the

electrophilic Cr(CO)3 fragment withdraws electon density from the aromatic ring activating it towards nucleophilic

Semmelhack JACS 1980 (102) 5926

π-acceptor effects on reactivity

electrophilic Cr(CO)3 fragment withdraws electon density from the aromatic ring activating it towards nucleophilic

attack.

Yamamoto JACS 1971 (93)3350.

Acrolein is thought to act as a π-acid, withdrawing electron density from the Ni(II) complex via π-backbonding and

promoting elimination of the diethyl fragment to reduce the metal.6

Dewar-Chatt-Duncanson model

Explain the trend in CO vibrations for the following series of compoundsExplain the trend in CO vibrations for the following series of compounds

v(CO) cm-1

[Ti(CO)6]2- 1748

[V(CO)6]- 1859

Cr(CO)6 2000

[Mn(CO)6 ]+ 2100

[Fe(CO)6 ]2+ 2204

7

The increase in electron density at the nickel from PR3 σ-

donation is dispersed through the M-L π system via π-

backbonding. Much of the electron density is passed onto

M-CO ππππ-Backbonding and the trans-effect

CO stretching frequencies

measured for Ni(CO)3L where L are

PR3 ligands of different σ-donor

abilities. [v(CO) =2143 cm-1]

backbonding. Much of the electron density is passed onto

the CO π* and is reflected in decreased v(CO) stretching

frequencies which corresponds to weaker CO bonds.

Recall: Band position in IR is governed by :

1. force constant of the bond (f) and

2. individual masses of the atoms (Mx and My).

Stronger bonds have larger force constants than weaker bonds.

Tolman Chem. Rev. 1977 (77) 313 8

• In the L→M π-donor system filled p ligand orbitals exist slightly lower in energy

than the metal t2g set with which overlap occurs.

• The bonding t2g MO’s formed are again lower in energy than the initial metal t2g

set, however, the corresponding t2g* anti-bonding MO’s are lower in energy than

the eg* σ anti-bonding orbitals.

• Additionally, as the ligand t2g p orbitals are typically of lower energy than the

metal t2g set, the bonding t2g MO’s formed are predominantly ligand based, with

the corresponding t * anti-bonding MO’s being predominantly metal based.

MO description for L→M ππππ-donor system in an Oh complex

the corresponding t2g* anti-bonding MO’s being predominantly metal based.

• Thus the LFSE ∆o is decreased relative to the σ only system. The t2g bonding MO’s

are stabilized which is countered by occupation of the t2g* anti-bonding orbitals.

• Overall this combined σ and π donation from ligand to metal results in an

increased M-L bond order and a stronger bond, however, the metal now becomes

more electron rich which can decrease the bond strengths of the remaining ligand

set making ligand-metal π bonding a less favorable interaction.

9

4s (a1g)

4p (t1u)

CrIII 6Cl-

eg*

t1u*

a1g*

Cl

CrCl ClCl

Cl

Cl

t2*

3-

T2g3 x

a1g

3d (t2g , eg)

t1u

eg

a1g

t1u

eg

t2g

t2g

t2g

dx2-y2dz2

10

MO description for L→M ππππ-donor system in an Oh complex

The energy of the HOMO is directly affected by M-L π-bonding. Ligand to metal π-

donation increases the energy of the HOMO making the metal more basic. π-donor

ligands stabilize electron poor, high oxidation state metals. Very prevalent for early TM

complexes (low d electron count) and less so for late TM (high d electron count).

11

• Alkoxy, amido and halido ligands are

sp2 hybridized so only one set of p

electrons are available for π bonding

12

Summary of ππππ-bonding in a Oh complexes

ππππ donor ligands result in L→M ππππbonding, a smaller ∆∆∆∆o favoring high spin

configurations and a decreased stability.

ππππ acceptor ligands result in M→L ππππ bonding, a larger ∆∆∆∆o

favoring low spin configurations with an increased

stability.13

Types of bonding

σ π δ σ π δ σ π δ σ π δ

• Bond polarity is evaluated by the difference in electronegativities of neighboring atoms

14

Ionic bonding

• Greater when elements of high and opposite charge interact.

• Differences in charge are paralleled in differences in

electronegativities.

• Large differences in electronegativity favor strong ionic bonding.

• M-L σ-bonding in early metals has significant ionic character.

M-L σ bond polarization

Covalent bonding

• Greater when orbitals of similar energies interact.

• The energy of atomic orbitals is inversely proportional to the

element's electronegativity (i.e. the orbital energy of an

electronegative element is lower than that of a electropositive

element).

• Small differences in electronegativity favor strong covalent bonding.

• M-L σ-bonding in late metals has a high degree of covalent bonding.

15

• Both the covalent model and the ionic model differ only in the way the electrons

are considered as coming from the metal or from the ligands

- emphasize model…not a true representation of metal charge!!!

• Each model is often invoked without any warning in the literature therefore it is

important to be able to identify their use.

• The ionic model is most commonly used for traditional M−−−−L inorganic

coordination compounds therefore coordinating ligands are treated equally in

both models.

Ionic vs. covalent model

both models.

• The ionic model is more appropriate for high-valent metals with N, O or Cl ligands.

• In the ionic model the M−X bond is considered as arising from a cationic M+ and

an anionic X− (heterolytic)

• The covalent model is sometimes preferred for organometallic species with low-

valent metals where the metal and ligand oxidation states cannot be

unambiguously defined.

• In the covalent model the M−X bond is considered as arising from a neutral metal

and ligand radical X• (homolytic).16

Hydride

Alkyl (e.g. methyl)

Alkenyl

Alkynyl

Allyl

-1

-1

-1

-1

-1

-1

2

2

2

2

2

2

0

0

0

0

0

0

1

1

1

1

1

1

ηηηη1 (hapticity = 1)

Ionic model covalent model

# of e- # of e-

charge donated charge donated

Aryl

Cyclopentadienyl

Carbene

Carbyne

Acyl

Carbon monoxide

Nitrile

Fulminate

-1

-1

-2

-3

-1

0

-1

-1

2

2

4

6

2

2

2

2

0

0

0

0

0

0

0

0

1

1

2

3

1

1

1

1

17

Aqua

Hydroxyl

Oxo

Alkoxide

Ether

(Sulfide/Tioether O = S)



# of e- # of e-


0

-1

-2

-1

0

-1

2

2

4

2

2

2

+1

0

0

0

+1

0

1

1

2

1

1

1

Carboxylate

Carbonate

Cyanate(Thiocyanate O = S)

Isocyanate(Isothiocyanate O = S)

Sulfoxide

-1

-1

-1

0

0

2

2

2

2

2

0

0

0

+1

+1

1

1

1

1

1

18

Isocyanate

Nitrogen

Amine

Pyridyl

Imine

Amide

Nitrile



# of e- # of e-


-1

0

0

0

0

-1

0

2

2

2

2

2

2

2

0

+1

+1

+1

+1

0

+1

1

1

1

1

1

1

1Nitrile

Isonitrile

Nitrosyls

Nitro

Nitrito

Phosphine

Phosphide

Halide (e.g. Cl)

0

0

-1

0

-1

-1

0

-1

-1

2

2

2

2

2

2

2

2

2

+1

+1

0

+1

0

0

+1

0

0

1

1

1

1

1

1

1

1

1 19

µµµµ-hydride

µµµµ-oxo

µµµµ-alkoxide

µµµµ-halide

µµµµ-CO

HM M

O

M M

O

M M

R

X

M M

M

C

M

O

H2N M'

Bridging ligands (non-chelating)


# of e- # of e-


n/a

-2

-1

-1

-2

0

n/a

4

4

4

4

2x2

-1

0

+1

-1

0

+2

2

2

2

2

2

2x1

µµµµ-en (ethylene diamine)

µµµµ-pyrazine

µµµµ-4,4’-bipyridine

µµµµ-dppe

[1,2-bis(diphenyl phosphino)ethane]

NH2M

NM N M

NM N M'

0

0

0

0

2x2

2x2

2x2

2x2

+2

+2

+2

+2

2x1

2x1

2x1

2x1

20

• Common unsaturated ππππ donating ligands encountered in organotransition-metal chemistry

together with the respective numbers of electrons relevant to the application of the 18 VE rule.21

• When applying the 18VE rule the following should be considered

1. The intramolecular partitioning of the electrons has to ensure that the total

charge of the complex remains unchanged (ionic or covalent model).

The 18 VE rule

2. A M−M bond contributes one electron to the total electron count of a single

metal atom.

3. The electron pair of a bridging ligand donates one electron to each of the

bridged metal atoms.

22

1. The intramolecular partitioning of the electrons has to ensure that the total

charge of the complex remains unchanged.

Electron counting - revision

23

2. A M−−−−M bond contributes one electron to the total electron count of a single

metal atom.

What is the d electron count of Mn in the unstable (CO)5Mn monomer?

24

3. The electron pair of a bridging ligand donates one electron to each of the

bridged metal atoms.

25

• Transition metal complexes can be divided into 3 classes:

Class # valence electrons 18 VE rule

I . . . 16 17 18 19 . . . not obeyed

II . . . 16 17 18 not exceeded

III 18 obeyed

The 18 VE rule

• How does the nature of the central atom and of the ligands determine whether a

complex belongs to class I, II, or III?

• Guiding principles to construction of a transition metal complex:

� antibonding orbitals should not be occupied

� nonbonding orbitals may be occupied

� bonding orbitals should be occupied

26

Class I

� The splitting of ∆∆∆∆0 is relatively small for 3d metals as well as for σ ligands at

the lower end of the spectrochemical series.

� t2g is non-bonding and can be occupied by 0 − 6 electrons.

� eg* is weakly anti-bonding and can be occupied by 0 − 4 electrons.� eg* is weakly anti-bonding and can be occupied by 0 − 4 electrons.

� Thus 12 −−−− 22 valence electrons can be accommodated, i.e. the 18VE rule is not

obeyed.

� Owing to their inherently small ∆0 splitting tetrahedral complexes also belong

to this class.

27

Class II

� ∆∆∆∆0 is larger for 4d and 5d metals (especially in the higher oxidation states) as

well as for strong σ donating ligands.

� t2g is essentially non-bonding and can be occupied by 0 − 6 electrons.

� eg* is more strongly anti-bonding and is no longer available for occupancy.� eg* is more strongly anti-bonding and is no longer available for occupancy.

� Consequently the valence shell contains 18 VE or less.

� A similar splitting ∆0 is also observed for complexes of 3d metals with ligands

that exhibit extremely high ligand-field strength (e.g. CN−)

29

Class III

� ∆0 is largest for ligands at the upper end of the spectrochemical series (good π

acceptors, such as CO , PF3, olefins, arenes).

� t2g is now bonding due to π interactions with orbitals of the ligands and

should be occupied by 6 electrons.

� eg* is strongly anti-bonding and remains unoccupied.

� The valence shell contains 18 VE obeying the 18VE rule.

� Exceptions can occur usually for steric reasons, e.g. V(CO)6 contains 17VE but

is readily reduced to [V(CO)6]−.

Organometallic complexes of transition metals almost exclusively belong to Class III

31

T2g3 x

dx2-y2dz2

33

• For M(d8) late transition metals 16 VE configurations are favored

Examples: M (d8)

[Ni(CN)4]2− 16 VE

[Rh(CO)2Cl2]− 16 VE square planar, CN 4

[AuCl4]− 16 VE

Deviations from the 18VE rule

• For M(d10) late transition metals 14 VE configurations are favored

Examples: M (d10)

[Ag(CN)]− 14 VE

Ph3PAuCl 14 VE linear, CN 2

• d8 (16e−−−−) complexes prefer square-planar

• d0, d5, d10 complexes usually tetrahedral34

• According to the 18VE rule, a coordination number (CN) of 5 is expected for d8

late transition metals.

• This is found for early transition metal complexes with a d8 configuration.

Examples: M (d8)

Fe(CO)5 18 VE

Mn(CO) −−−− 18 VE trigonal bipyramidal, C.N. 5Mn(CO)5−−−− 18 VE trigonal bipyramidal, C.N. 5

[(ηηηη4−−−−C4H8)(ηηηη6−−−−C6H6 )Ru] 18 VE

35

• A qualitative explanation for deviations from the 18VE rule takes into a/c the

following:

� electroneutrality principle (Pauling, 1948)

� ππππ−−−−donating character of the ligands

� change in energy separation between (n – 1)d, ns, and np orbitals within a � change in energy separation between (n – 1)d, ns, and np orbitals within a

transition metal series.

36

• Increase in atomic number is accompanied by a decrease in orbital energies

across the periodic table as additional valence electrons only partially screen the

higher nuclear charge.

� For (n-1)d electrons, the decrease in energy is more pronounced than for ns

and np electrons

(d orbitals are more diffuse and experience a lower nuclear shielding relative (d orbitals are more diffuse and experience a lower nuclear shielding relative

to s and p orbitals of the same principle quantum number)

� Successive addition of d electrons across a transition series therefore

increases the effective nuclear charge and stabilizes the d orbitals.

� The energy separation between the (n-1)d, ns, np atomic orbitals also

increases with an increasing positive charge of the atom.

37

Electroneutrality Principle

“Stable complexes are those with structures such that each atom has only a small

electric charge. Stable M-L bond formation generally reduces the positive charge

on the metal as well as the negative charge and/or e- density on the ligand. The

result is that the actual charge on the metal is not accurately reflected in its

formal oxidation state”

- Pauling; The Nature of the Chemical Bond, 3rd Ed.;1960, pg. 172.- Pauling; The Nature of the Chemical Bond, 3rd Ed.;1960, pg. 172.

• The 18VE rule often fails for early transition metals.

• Formal oxidation state is not an accurate description of electron density at the

metal.

• High oxidation state, early TM complexes are stabilized via π-donation

i.e. a shifting of electron density from π-donor ligands to the metal.

• This in part accounts for the extreme oxophilicity of early TM.

38

1) Determine the oxidation state of the metal

(Balance ligand charges with overall charge

of the complex)

2) Determine the number of valence electrons

d electron count = (Group #) – (metal

oxidation state)

Revision: Electron Counting

3) Count electrons donated by neutral

or anionic ligands

(bridging ligands donate half electrons to

either metal per coordination site)

4) Metal – Metal bonds

(one electron per bond)

39

Homework: Construct a ππππ MO diagrom for Titanium

Tetraisopropoxide, Ti(OiPr)4

41

General properties of organometallic complexes

• There are several differences between organometallic complexes and coordination

compounds:

� The metals are more electron rich, i.e. the metal bears a greater negative charge in the

organometallic complex.

� The M−L bonds are much more covalent and often have a substantial π component.

� The metal d orbitals are higher in energy and by back donation perturb the electronic

structure of the ligands much more than is the case for coordination compounds.

� The organometallic ligands can be polarized and therefore activated toward chemical

reactions, σ and π bonds in the ligands can be weakened or broken, and chemical bonds

can be made or broken within and between different ligands.

� Readiness to change the coordination number and the lability of the M−C σ bond are

essential for organometallic catalysis

• This rich pattern of reactions is characteristic of organometallic chemistry.

42

valence bond description of the co ligand

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