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.

δ− δ+ • This transfer leads to a C −−−O p p polarization of the molecule, which is almost exactly canceled out by a partial Cδ+−−−Oδ− polarization of all three sp C sp sp O sp bonding orbitals because of the higher p dative p electronegativity of oxygen. • The free CO molecule therefore has a net dipole moment very close to zero.

1 MO correlation diagram for CO

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 C metal dxy, dxz and dyz orbitals. C O O

LUMO

HOMO 2p

2p

The CO HOMO orbital is a bonding orbital of σ symmetry with significant electron density on the 2s carbon. This orbital forms a σ bond with metal p , dz2 and dx2–y2 orbitals. This is a filled orbital and 2s donates electron density to the metal centre.

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

2 3 x T2g

LUMO

∆∆∆O

HOMO dz2 dx2-y2

3 4 • 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.

2−
If the L groups are good donors or the complex is anionic, e.g. Cp 2W(CO) or [W(CO) 5] , 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 σ ∗ ∗ donor, through to the extreme in which both the π x and π y are both fully engaged in back bonding.

5 π-acceptor effects on reactivity

Semmelhack JACS 1980 (102) 5926 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 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 compounds v(CO) cm -1

2- [Ti(CO) 6] 1748 - [V(CO) 6] 1859

Cr(CO) 6 2000 + [Mn(CO) 6 ] 2100 2+ [Fe(CO) 6 ] 2204

7 M-CO πππ-Backbonding and the trans-effect

The increase in electron density at the nickel from PR 3 σ- donation is dispersed through the M-L π system via π- 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.

CO stretching frequencies Recall: Band position in IR is governed by : measured for Ni(CO) 3L where L are 1. force constant of the bond ( f) and PR ligands of different σ-donor 3 2. individual masses of the atoms (M and M ). abilities. [v(CO) =2143 cm -1] x y Stronger bonds have larger force constants than weaker bonds. Tolman Chem. Rev. 1977 (77) 313 8 MO description for L→M πππ-donor system in an Oh complex

• 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 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 3- Cl CrIII 6Cl- Cl Cl Cr Cl Cl Cl

* t1u

3 x T2g

4p (t1u) * a1g

4s (a1g)

* eg * t2g

3d (t2g , eg)

dz2 dx2-y2 t2g

t2g eg

t1u eg

a1g t1u

10 a1g MO description for L→M πππ-donor system in an O h 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 sp 2 hybridized so only one set of p electrons are available for π bonding

12 Summary of πππ-bonding in a O h complexes

πππ donor ligands result in L→M πππ bonding, a smaller ∆∆∆o favoring high spin configurations and a decreased stability.

πππ acceptor ligands result in M→L πππ bonding, a larger ∆∆∆o favoring low spin configurations with an increased stability. 13 Types of bonding

σ π δ

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

14 M-L σ bond polarization

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.

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 Ionic vs. covalent model

• 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 . • 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 Ionic model covalent model ηηη1 (hapticity = 1) # of e- # of e- charge donated charge donated

-1 2 0 1 Hydride -1 2 0 1 Alkyl (e.g. methyl)

-1 2 0 1 Alkenyl

-1 2 0 1 Alkynyl

-1 2 0 1 Allyl

-1 2 0 1 Aryl

-1 2 0 1 Cyclopentadienyl

-2 4 0 2 Carbene

-3 6 0 3 Carbyne

Acyl -1 2 0 1

0 2 0 1 Carbon monoxide -1 2 0 1 Nitrile -1 2 0 1 Fulminate 17 Ionic model covalent model ηηη1 (hapticity = 1) # of e- # of e- charge donated charge donated

0 2 +1 1 Aqua

-1 2 0 1 Hydroxyl -2 4 0 2 Oxo -1 2 0 1 Alkoxide

0 2 +1 1 Ether (Sulfide/Tioether O = S)

-1 2 0 1 Carboxylate

-1 2 0 1 Carbonate

-1 2 0 1 Cyanate (Thiocyanate O = S) Isocyanate (Isothiocyanate O = S) -1 2 0 1

Sulfoxide 0 2 +1 1

0 2 +1 1

18 Ionic model covalent model ηηη1 (hapticity = 1) # of e- # of e- charge donated charge donated Isocyanate -1 2 0 1 Nitrogen 0 2 +1 1 Amine 0 2 +1 1

0 2 +1 1 Pyridyl

Imine 0 2 +1 1

Amide -1 2 0 1 Nitrile 0 2 +1 1

0 2 +1 1 Isonitrile -1 2 0 1 Nitrosyls

0 2 +1 1

-1 2 0 1 Nitro

Nitrito -1 2 0 1

0 2 +1 1 Phosphine -1 2 0 1 Phosphide -1 2 0 1 Halide (e.g. Cl) 19 Ionic model covalent model Bridging ligands (non-chelating) # of e- # of e- charge donated charge donated

M H M µµµ-hydride n/a n/a -1 2 O µµµ-oxo M M -2 4 0 2 R

O -1 4 +1 2 µµµ-alkoxide M M X M M -1 4 -1 2 µµµ-halide O

C M M -2 4 0 2 µµµ-CO

H2N M' 0 2x2 +2 2x1 µµµ-en (ethylene diamine) M NH2

µµµ-pyrazine M N N M 0 2x2 +2 2x1

µµµ-4,4’-bipyridine M N N M' 0 2x2 +2 2x1

µµµ-dppe [1,2-bis(diphenyl phosphino)ethane] 0 2x2 +2 2x1

20 • Common unsaturated πππ donating ligands encountered in organotransition-metal chemistry

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

• 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).

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 Electron counting - revision

1. The intramolecular partitioning of the electrons has to ensure that the total charge of the complex remains unchanged.

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 The 18 VE rule

• 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

• 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.
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 28 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.
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 30 Class III

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

acceptors, such as CO , PF 3, 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 32 3 x T2g

dz2 dx2-y2

33 Deviations from the 18VE rule

• For M(d 8) late transition metals 16 VE configurations are favored Examples : M (d 8)

2− [Ni(CN) 4] 16 VE

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

− [AuCl 4] 16 VE

• For M(d 10 ) late transition metals 14 VE configurations are favored Examples : M (d 10 ) [Ag(CN)] − 14 VE

Ph 3PAuCl 14 VE linear, CN 2

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

0 5 10 • d , d , d complexes usually tetrahedral 34 • According to the 18VE rule, a coordination number (CN) of 5 is expected for d 8 late transition metals.

• This is found for early transition metal complexes with a d 8 configuration.

Examples : M (d 8)

Fe(CO) 5 18 VE

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

4 6 [( ηηη −−−C4H8)( ηηη −−−C6H6 )Ru] 18 VE

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

electroneutrality principle (Pauling, 1948)
πππ−π−−−donating character of the ligands
change in energy separation between (n – 1)d , n s, and n p 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 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.

• 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 Revision: Electron Counting

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)

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 40 Homework: Construct a πππ MO diagrom for Titanium i Tetraisopropoxide, Ti(O Pr)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