Inorganic Chemistry - Metal Complex Ligand

Organometallic Chemistry

• Introduction to Organometallic Chemistry.
• Explanation of π-donor and π-acceptor ligands using Ligand Field Theory.
• Dewar-Chatt-Duncanson Theory.
• The 18 electron rule.

The Birth of Organometallic Chemistry and Carbonyl Compounds

• Louis Claude Cadet de Gassicourt (1731-1799) is credited with the first recorded organometallic compounds in the 1760s in France.
• Cadet prepared dicacodyl (As2Me4) from arsenic salts, a repulsive and toxic substance that fumes in the air and smells like garlic.

Metal-Carbon Bonds

• Metal alkyl and aryl (M-C bonds) are formed in organometallic chemistry.
• Examples include Ferrocene and Zeise’s salt.
• François Grignard was awarded the 1912 Nobel Prize in Chemistry for his work in this area.
• Transition metals facilitate diverse organometallic compounds due to available d-orbital geometries for coordinate bonding with pi orbitals of unsaturated molecules by donating and accepting electrons.

The 18 Electron Rule

• Definition: A stable organometallic complex results when the sum of the metal valence electrons plus the total number of electrons donated by the ligands equals 18.
• This rule is the transition metal (TM) equivalent of the octet rule for p-block elements.
• Transition metals have 5 extra valence orbitals, allowing for 10 extra valence electrons in bonding molecular orbitals.
• Early transition metals commonly exhibit 16e and sub-16e configurations, often with coordination numbers greater than 6.
• Middle transition metals commonly have 18e configurations with coordination numbers equal to 6.
• Late transition metals commonly show 16e and sub-16e configurations, frequently with coordination numbers of 5 or less.

Exceptions to the 18 Electron Rule

• The 18 electron rule is a useful predictive tool but is not absolute; many complexes do not follow it.
• Exceptions occur especially when:
  • The splitting energy (∆) is small.
  • Steric factors are significant, e.g., V(CO)6, Cr(1-adamantyl)4.
  • The metal has a low electron count (early transition metals), e.g., WMe6, TiCl4(thf)2.
  • Electron-rich d8 metal complexes are square planar, possess 8 low-lying molecular orbitals and abide by a 16 electron rule!

Electron Counting

• Steps for electron counting using the 18 electron rule:
  • Count the total number of electrons donated by ligands.
  • Calculate the oxidation state (O.S.) of the metal center:
    • From charges on ligands.
    • From overall charge on the complex: O.S. = (overall charge) – (total charge from the ligands)
  • Determine the number of d electrons: Number of d electrons = Element group number – O.S.
  • Calculate total valence electrons: Total VE = No. of d electrons from metal + No. of electrons donated by ligands

Electron Counting Examples

• Example 1:
  • Total electrons from ligands = 6 x 2 = 12
  • Oxidation state of metal = 0 – (0) = 0
  • Number of metal d electrons = 6
  • Total valence electrons = 18
• Example 2:
  • Total electrons from ligands = 4 x 2 = 8
  • Oxidation state of metal = 0 – (-1) = 1
  • Number of metal d electrons = 8
  • Total valence electrons = 16
• Example 3:
  • Total electrons from ligands = 8 x 2 = 16
  • Oxidation state of metal = -2 – (-8) = 6
  • Number of metal d electrons = 1
  • Total valence electrons = 17

Sigma and Pi Bonding in Complexes

• Crystal Field Theory (CFT) previously considered ligands as “ionic interactions.”
• In reality, covalent character and orbital overlap exist.
• Further investigation of bonding requires a combination of CFT and molecular orbital theory.
• σ- and π-type bonding in coordinate bonds must be considered.

Sigma-Bonding (σ)

• Overlap occurs for the 4s, 4p, and eg orbitals with the ligands in pure sigma bonds.
• The t2g orbitals are non-bonding because they do not overlap along the bonding axis.

Pi-Acid (Acceptor) Ligands

• π-acceptor ligands accept electron density from metals through π-type (perpendicular) bonding.
• To act as a π-acid, the ligand must have an empty orbital of similar energy and symmetry.
• Molecular orbital diagram for CO:
  • HOMO (Highest Occupied Molecular Orbital) is Sσ* orbital.
  • LUMO (Lowest Unoccupied Molecular Orbital) is a pπ* degenerate pair.

Pi-Back-Bonding

• Two major components in transition metal carbonyl complexes:
  1. σ-donation from Sσ* orbital of CO.
  2. π-donation from a metal d-orbital to the CO (pπ*).
• The latter is referred to as “back-bonding,” and the overall situation is defined as synergic, i.e., mutually reinforcing.
• σ-donation Stabilizes C≡O bond by removing electron density from the anti-bonding orbital.
• π-donation Destabilizes C≡O bond by adding electrons into the anti-bonding orbital.

Synergic Bonding

• IR spectroscopy is the most sensitive probe for CO bond strength.
• CO strength decreases as electron density on the metal (and hence π back-donation) increases.
• σ-donation depletes electron density from the CO anti-bonding orbital, increasing the C-O bond strength.
• π-donation places electron density into anti-bonding ligand orbitals, reducing C-O bond order and decreasing CO bond strength.
• CO is a powerful π–acceptor but a poor σ-donor.
• CO rarely coordinates to metals unable to back-donate.
• In the absence of any coordination, C≡O has a strong IR stretching frequency (2143 cm-1).

Examples of Synergic Bonding - IR

• Complex [Mn(CO)6]+ has ν(CO) at 2090 cm-1.
• Complex Cr(CO)6 has ν(CO) at 2000 cm-1.
• Complex [V(CO)6]- has ν(CO) at 1860 cm-1.
• Complex Ni(CO)4 has ν(CO) at 2060 cm-1 (Tetrahedral).
• Complex [Co(CO)4]- has ν(CO) at 1890 cm-1.
• Complex [Fe(CO)4]2- has ν(CO) at 1790 cm-1 (Octahedral).
• The stretching frequency decreases, indicating a lower CO bond order due to increased back-donation.

The Nitrosyl (NO) Ligand

• The bonding in the NO radical resembles CO except that NO has 1 extra electron occupying one pπ* orbital; the N-O bond order = 2.5.
• The loss of this electron gives NO+, which is isoelectronic with CO but is a cationic 2e- donor.
• The relevant Sσ and pπ orbitals are lower in energy in NO+, compared to CO, making it a poorer σ-donor but stronger π-acceptor.
• NO can also add an electron to give NO-, an anionic 2e- donor, with a formal N=O bond.
• The nature of the NO donor can be determined if the M-N-O bond angle is known.
• NO+ is ≈180 (linear), and NO- is ≈120 (Bent).

Electron-Counting in NO complexes

• Example 1 (all four ligands are NO+):
  • Total electrons from ligands = 4 x 2 = 8
  • Oxidation state of metal = 0 – (4) = -4
  • Number of metal d electrons = 10
  • Total valence electrons = 18 (Obeys 18e- rule)
• Example 2 (all four ligands are NO-):
  • Total electrons from ligands = 4 x 2 = 8
  • Oxidation state of metal = 0 – (-4) = 4
  • Number of metal d electrons = 2
  • Total valence electrons = 10

Pi-Base (Donor) Ligands

• Any ligand can act as a π-base through donation of a second lone-pair to an appropriate, empty metal orbital, forming a formal M=L double bond.
• Common examples are amido (R2N-) and alkoxide (RO-) ligands that have filled p-orbitals which can form molecular orbitals via a π-bond with d-orbitals.
• π-donor complexes result in a decrease in the apparent ligand field (Δ).

Pi-Base (Donor) Ligands

• π-donating ability of a ligand decreases as the overall charge on the ligand decreases and/or electronegativity of the bonding atom increases.
• Comparison of π-donor ability: RC3- > N3- > R2C2- > RN2- > O2- > R2N- > RO-

Evidence for Pi-Bases and Pi-Acids

• Large Δ, strong-field ligands = π-acceptor ligands
• Small Δ, weak-field ligands = π-donor ligands
• Also inferred from bond lengths obtained from x-ray crystallography.

Pi-Effects in Octahedral complexes

• Ligand-to-Metal Charge Transfer (LMCT) observed for π-donor ligands.
• Metal-to-Ligand Charge Transfer (MLCT) observed for π-acceptor ligands.

The Dewar-Chatt-Duncanson Model

• Dewar-Chatt-Duncanson model for bonding in a TM-alkene complex:
  1. σ–donation of a π-bonding electron-pair on the alkene to an empty σ-type metal orbital
  2. π–donation from a filled metal d-orbital into an empty π*-orbital on the alkene.
• σ-bonding: σ-component: C-C p → empty metal orbital.
• π-back bonding: p-component: occupied metal d → empty C-C p*.

Bonding in TM-Alkene Complexes

• Like CO complexes, the σ-component of the M-C bond is weak, with π-back-bonding interactions dominating.
• Both the σ-donation and π-acceptance increase the M-C bond strength (synergic) and weaken the C=C bond of the alkene.

Bonding in TM-Alkene Complexes

• Weak π-basic metals (i.e., Pd(II), Pt(II)):
  • Alkene acts as a neutral 2e donor.
  • Alkene geometry is weakly σ-bonded with little π-back-donation.
  • Example: [PtCl3(C2H4)]−.
• Strong π-basic metals (i.e., Pd(0), Pt(0)):
  • Form metallacyclopropane structure.
  • Alkene acts as a dianionic 4e donor.
  • sp³ carbons.
  • Strong σ-donation from two lone pairs, no π-contribution.
  • Example: (Cp*)2Zr(C2H4), Fe(CO)4(C2H4).

Evidencing Bonding in TM-Alkene Complexes

• Alkenes can rotate about the M-C=C axis:
  • M-alkene σ-bond is invariant.
  • M-alkene π-bond is STRONGLY DEPENDENT on rotation angle.
  • π-backbonding requires a specific, coplanar orbital interaction.
• M-alkene bonding with an appreciable π-component restricts ligand rotation around the metal.
• This can be observed by variable temperature NMR.

Phosphine Ligands

• Phosphines, PR3, are neutral 2e- donors.
• Donor properties (electronic and steric) can be easily modified through changes in their R group.
• The combination of phosphorus d orbitals and a degenerate set of σ*-orbitals leads to the formation of π-type receptor orbitals.
• Phosphines are generally strong sigma-donors that can also act as π-acids.
• The extent of σ-donation and π-acceptance is dependent on the R group.

Phosphine Ligands

• σ-donation is enhanced by electron-donating groups.
• π-acceptance is enhanced by electron-withdrawing groups.
• Greater σ-basicity/ π-acceptance: PMe3 ≈ P(NR2)3 < PPh3 < P(OMe)3< PCl3 < PF3 ≈ CO
• M-P bond distance decreases (strengthens) upon reduction of the metal center.
• The P-R bond lengthens (weakens). The extent of this depends on R.

Phosphine Ligands

• Tolman, 1972 – Measured σ-donor strength by comparing C-O stretching frequency in FT-IR spectroscopy.

Phosphines

• Show a preference for later, electron-rich transition metals, especially when good π-acids.
• Only strongly σ-basic phosphines form stable complexes with high-valent early transition metals.

Summary

• Availability of d-orbitals allows diverse bonding modes and chemistry for organometallic chemistry.
• Spectrochemical series of complexes can be explained by covalent-like interaction of the ligands and the central metal atom.

Learning Outcomes

• Understand and explain the differences between π-donor and π-acceptor ligands using Ligand Field Theory
• Explain the Dewar-Chatt-Duncanson theory of Metal-Ligand bonding
• Be able to count electrons in transition metal complexes, using the 18 e- rule

Revision Topics

• [Mn(H2O)6]3+ has an μeff of 4.9 BM. How many unpaired electrons does it have and is it a high or low spin complex?
• Determine the oxidation state, the d configuration, the number of unpaired electrons, and whether the complex is high or low spin for:
  • [Co(H2O)6]3+ 4.9 BM
  • [Co(CN)6]3 - 0 BM
  • [Co(NO2)6]4 - 1.8 BM
• Answers:
  • [Co(H2O)6]3+ 4.9 BM (n=4); Co3+ d6 high spin
  • [Co(CN)6]3 - 0 BM (n=0 ); Co3+ d6 low spin
  • [Co(NO2)6]4 - 1.8 BM (n=1); Co2+ d7 low spin
• [Fe(CN)6]4 - Δo = 33000 cm-1 and [Fe(H2O)6]2+ Δo= 10400 cm-1 has pairing energy is 17600 cm-1, determine the electron configurations for each complex and Identify the type of magnetism demonstrated by each complex.
• [CoF6]3 - , with 6 ligands gives Oh geometry, Δo = 13,000 cm-1. Co3+ is d6 with 4 unpaired electrons and so is a paramagnetic, high spin, complex. High spin complex, therefore, spin pairing energy, P, must be greater than 13,000 cm-1.