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Classification of metal Carbonyls
(I) On the Basis of Ligands: Metal carbonyls can be classified into two categories
(a) Homoleptic carbonyl complexes: The complexes in which a metal is bound to only CO ligands are known as homoleptic carbonyl complexes. For example, Ni(CO)4, Cr(CO)6, Fe(CO)5 etc.
(b) Heteroleptic carbonyl complexes: The complexes in which a metal is bound to CO as well as other ligand such as PR3, PP3, NO etc. For example Ni(CO)3(PPh3), Mo(CO)3(PF3)3, etc.
(II) On the basis of number of metal atoms and the Structures of metal carbonyls:
(a) Mononuclear Metal Carbonyls: These carbonyls contain only one metallic atom and these carbonyls do not contain any bridging CO ligand.
(b) Polynuclear Metal Carbonyls: Polynuclear carbonyls contain two or more metal atoms.
Bonding in Metal Carbonyls CO is considered to be a good σ-donor as well as a π-acceptor ligand. This property of CO and the other π-acceptor ligands such as CN-, NO, C2H4 etc, can be explained b y the MO diagram of CO.
When energy difference between 2s and 2p-orbitals is small there will be mixing of s and pz orbitals. The lower energy atomic orbitals of oxygen contribute more to the bonding molecular orbital and the higher energy atomic orbital of carbon contribute more to anti bonding molecular orbitals. Therefore, in CO the bonding molecular orbitals will have the character of orbitals of oxygen and the anti bonding molecular orbitals have the character of orbitals of less electronegative carbon. This is due to the conversion of orbitals.
The molecular orbitals energy level diagram shows that the HOMO has σ symmetr y and the HOMO is localized on carbon, not on oxygen because the effective nuclear charge or electronegativity of carbon is less than that of oxygen. The HOMO of CO ligand donates its lone pair of electrons to the empty orbital of suitable symmetry on the metal to form a M—CO σ-bond.
Carbon monoxide has two LUMO π* orbitals which are also localized on carbon. These orbitals have correct symmetry to overlap with non-bonding metal d-orbitals that have π symmetr y such as the t2g (dxy, dyz or dzx) orbitals in octahedral complex. A metal atom having electrons in a d-orbital of suitable symmetr y can donate electron density to the LUMO π* of CO. The π interaction leads to the delocalization of electrons from filled d-orbitals of suitable symmetry on the metal atom into the empty π* orbitals on the CO ligands, this donation is known as π back donation or back bonding and the CO ligand is said to be a strong π-acceptor.
Both the σ - and π-bonding reinforce each other. The formation of s-bond results in the increase in electron density on metal and tends to make the CO ligand positive. Both the increase in electron density on the metal, and the positive charge on CO increase the p-accepting ability of CO. The greater electron density on the metal and greater partial positive charge on CO ligand make the effective return of electron densit y from metal d-orbital to the π* orbital of CO ligand. As the delocalized electron density increases from metal d-orbital to π* orbitals of CO, further electron flow from CO to metal dorbital results. The result of this two way electron flow is that the metal-ligand bond is stronger than the sum of isolated to metal s-bonding and metal to ligand p-bonding effects. This kind of mutual strengthening of σ-and π-bonding is called synergism and this effect is called synergistic effect.
The metal to ligand p back bonding results in the increase in M—C bond strength as the bond order tends to increase. Since the p-back bonding results in occupation of π* on CO ligand, the bond order of CO ligand itself decreases and, therefore, the C—O bond becomes and weaker. Alternatively, increased M—C double bonding leads to decreases C—O multiple bonding as shown by the resonance structure.
If a metal is in high oxidation state, then there will be poor p-back bonding and the bond order of M—C is close to 1 and that of C—O is close to 3.
M—C ≡ 0
If the metal is in low oxidation state, then there will be strong p-bonding between metal and CO and the bond order of both M—C and C—O is close to 2.
M—C = 0
Factors Affecting the Magnitude of Stretching Frequency
IR stretching frequency ∝ Bond order
(1) Charge on Metal: Greater the negative charge on metal lesser will be the stretching frequency of C—O bond and greater will be the stretching frequency of M—C bond.
Greater the positive charge on metal, greater will be the stretching frequency of C—O bond and lesser will be the stretching frequency of M—C bond. (For isoelectronic species generally)
(2) Number of CO Ligands: Generally greater the number of CO ligand around the central metal atom lesser will be stretching frequency of C—O bond due to increasing of formal negative charge on metal.
(3) Presence of the Other Ligand on Metal: Greater the donation tendency of electron density and lesser the accepting tendency of electron density lesser will be the stretching frequency of C—O bond and greater will be the stretching frequency of M—C bond due to increasing of formal negative charge on metal.
Bonding Modes of CO The Stretching frequency of bridging C—O bond is lesser than the non-bridging C—O bond because in bridging CO two metals donate its electron density to CO ligand while in non-bridging CO only one metal donates its electron density to CO ligand.
When the carbonyl ligand bridges three metals then its C—O bond stretching frequency becomes lesser than the C—O bond stretching frequency which is attached to two metals. Greater the bridging of CO ligand to metals lesser will be the stretching frequency of C—O bond.
|Bonding Modes of CO||νCO(in cm–1)|
Metal Carbonyl Clusters The metal carbonyl clusters are classified into two categories:
(1) Low nuclearity carbonyl clusters: The low nuclearity carbonyl clusters contain met al atoms
Illustration: CO2(CO)8: Total number of valance electrons = 2 × 9 + 2 × 8 = 34
Number of Co—Co bond =
i.e., One Co—Co bond per metal.
Illustration: Fe3(CO)12: Total number of valance electrons = 3 × 8 + 12 × 2 = 48
Number of Fe—Fe bond
Total number of Fe—F e bonds = 3 i.e., number of Fe—Fe bonds per Fe atom = 2
(2) High Nuclearity Carbonyl Clusters: High nuclearity carbonyl cl usters have metal atoms ≥ 5, each form atleast One M—M bond.
The Isolobal Analogy: Some of the reactions of the metal carbonyls are parallel with main group non-metals and compounds. For example, chlorine atom and methyl (CH3· ) free radical both have 7 valence electrons, one short of a noble gas configuration. Three of sp3 orbitals of carbon are involved in t he formation of s-bonds with the hydrogens. The fourth sp3 orbital is singly occupied and has higher energy than the bonding ones. The Mn(CO)5 fragment has 17 valence electrons, one short of the 18-electron. The σ -bonding between Mn and CO ligands in this fragment may be considered to involve the five of d2sp3 hybrid orbitals of Mn. The sixth hybrid orbital is singly occupied and has higher energy than the five σ -bonding orbitals. Some isolabal Frangments of Transition Metals and Main Group Elements
The ionic clusters of main group elements are called Zintl ions. Since these ions have no ligand, these are also known as nacked clusters.
e.g. Pb52- , Bi33- , Sn52- , Bi53+ , Te62+ , etc.
Zintl ions are also classified into closo, nido, arachno and hypho clusters.
If total number of valance electrons follows the pattern
|Type of total valance electrons system||Name of cluster|
|4n + 2||Closo|
|4n + 4||Nido|
|4n + 6||Arachno|
|4n + 8||Hypho|
Where n = number of metal atoms in the Zintl ion
Total number of valance electrons = 4 × 5 + 2 = 22
∴ n = 5
∵ type of total valance electron system = 4n + 2
i.e., Zintl ion has closo structure