Bonding in Metal Carbonyls: Metal carbonyls are coordination complexes of transition metals surrounded by carbon monoxide ligands. The first known metal carbonyl compound to be synthesized is \({\rm{Ni}}{\left( {{\rm{CO}}} \right)_4}\) by Ludwig Mond in \(1884\). The carbon monoxide ligand is a unique ligand and distinguishes itself from other ligands in many respects. Unlike the alkyl ligands, the carbonyl \({\rm{(CO)}}\) ligand is unsaturated, which means it not only donates electrons to form \(\sigma – \) bond but accepts electrons in its \({\pi ^*}\) orbital from \({{\rm{d}}_\pi }\) metal orbitals. This makes the \({\rm{CO}}\) ligand acidic. Moreover, \({\rm{CO}}\) is a soft ligand compared to other ligands like \({{\rm{H}}_2}{\rm{O}}\) or alkoxides \(\left( {{\rm{R}}{{\rm{O}}^ – }} \right)\). Let’s discuss how the \({\rm{CO}}\) ligand is bonded to metal carbonyls.

A lone pair of electrons is available on both carbon and oxygen atoms in the carbon monoxide ligand. However, as the carbon atoms donate electrons to the transition metals, the formed metal complexes are called carbonyls. In this article, let’s learn everything about Metal Carbonyls in detail.

Molecular Orbital Diagram of CO

To understand the bonding in metal carbonyls, we need to first learn the Molecular Orbital \(\left( {{\rm{MO}}} \right)\) diagram of carbon monoxide. There are ten electrons in the carbon monoxide ligand. The order of energy of the molecular orbitals and the accommodation of ten electrons of the carbon monoxide can be shown as:

The molecular orbital configuration of the \({\rm{CO}}\) molecule is :

\({\left( {\sigma {\rm{1}}\;{{\rm{s}}^{\rm{b}}}} \right)^{\rm{2}}}{\left( {\sigma {\rm{1}}\;{{\rm{s}}^ * }} \right)^{\rm{2}}}{\left( {\sigma {\rm{2}}\;{{\rm{s}}^{\rm{b}}}} \right)^{\rm{2}}}{\left( {\pi {\rm{2P}}{{\rm{x}}^{\rm{b}}}{\rm{ = }}\pi {\rm{2P}}{{\rm{y}}^{\rm{b}}}} \right)^{\rm{4}}}{\left( {\sigma {\rm{2P}}{{\rm{z}}^{\rm{b}}}} \right)^{\rm{2}}}{\left( {\sigma {\rm{2}}\;{{\rm{s}}^ * }} \right)^{\rm{2}}}\left( {\pi {\rm{2P}}{{\rm{x}}^ * }{\rm{ = }}} \right.{\left. {\pi {\rm{2P}}{{\rm{y}}^ * }} \right)^{\rm{0}}}{\left( {\sigma {\rm{2P}}{{\rm{z}}^{\rm{*}}}} \right)^{\rm{0}}}\)

In the molecular orbital of \({\rm{CO}}\)

1. The highest occupied molecular orbital (HOMO) that can donate the lone pair of electrons to form an \({\rm{OC}} \to {\rm{M}}{\mkern 1mu} \,{\rm{\sigma }}\) bond is \(\sigma 2\;{{\rm{s}}^*}\).

2. The lowest unoccupied molecular orbitals (LUMO) that can accept the electron density from an appropriately oriented filled metal orbital resulting in the formation of an \({\rm{M}} \to {\rm{CO}}\,{\rm{\pi }}\) bond is \(\pi {\rm{2P}}{{\rm{x}}^ * }{\rm{ = }}\pi {\rm{2P}}{{\rm{y}}^ * }\).

The carbon monoxide ligand is bonded to the transition metal through the formation of a dative \(\sigma \)-bond and \({\rm{\pi }}\)-bond due to back donation.

Formation of Dative σ-bond

The overlapping of the filled hybrid orbital (HOMO) on the carbon atom of carbon monoxide with an empty hybrid orbital on a metal atom results in the formation of an \({\rm{M}} \leftarrow {\rm{CO}}\,{\rm{\sigma }}\)-bond. The central metal in metal carbonyls is a transition element.

There is a vacant d-orbital available on the metal atom to accept the lone pair of electrons from the HOMO of the \({\rm{CO}}\) molecule. As we know, there are five d-orbitals, \({{\rm{d}}_{{\rm{xy}}}},{{\rm{d}}_{{\rm{yz}}}},{{\rm{d}}_{{z^2}}},\;{{\rm{d}}_{{\rm{zx}}}},{{\rm{d}}_{{{\rm{x}}^2} – {{\rm{y}}^2}}}\), available on the metal atom that can participate in the \({\rm{M}} \leftarrow {\rm{CO}}\,{\rm{\sigma }}\)-bond, the orbital with the proper alignment and filled orbital of \({\rm{CO}}\) is involved in the \({\rm{M}} \leftarrow {\rm{CO}}\,{\rm{\sigma }}\)-bond.

The donation of electrons takes place from the filled sigma molecular orbital, HOMO \(\left( {\sigma 2\;{{\rm{s}}^*}} \right)\) of the \({\rm{CO}}\) molecule. Hence, the d-orbital that has an alignment similar to the sigma orbital or HOMO of the \({\rm{CO}}\) molecule participates in the bonding.

From the image above, the \({{\rm{d}}_{{{\rm{x}}^{\rm{2}}}{\rm{ – }}{{\rm{y}}^{\rm{2}}}}}\) orbital has an alignment similar to that of the s orbital, HOMO of the \({\rm{CO}}\) molecule. The \({{\rm{d}}_{{{\rm{x}}^{\rm{2}}}{\rm{ – }}{{\rm{y}}^{\rm{2}}}}}\) orbital is along the XY axis, which is in alignment with the s orbital.

Hence, out of the five d orbitals, \({{\rm{d}}_{{{\rm{x}}^{\rm{2}}}{\rm{ – }}{{\rm{y}}^{\rm{2}}}}}\) participates in the \({\rm{M}} \leftarrow {\rm{CO}}\,{\rm{\sigma }}\)-bonding. The \({\rm{M}} \leftarrow {\rm{CO}}\,{\rm{\sigma }}\)-bond is diagrammatically represented as below-

Formation of π-back Bond

The π-bond in metal carbonyls is formed by overlapping filled dπ orbitals or hybrid dpπ orbitals of a metal atom with low-lying empty (LUMO) orbitals of \({\rm{CO}}\) molecule. i.e., \({\rm{M}} \to {\rm{CO\pi }}\) bonding.

The lowest unoccupied molecular orbitals (LUMO) of \({\rm{CO}}\) is \(\left( {{\rm{\pi 2P}}{{\rm{x}}^{\rm{}}}{\rm{ = \pi 2P}}{{\rm{y}}^{\rm{}}}} \right)\) which can accept the electron density from an appropriately oriented filled metal orbital. This results in the formation of an \({\rm{M}} \to {\rm{CO\pi }}\) bond. As the formation of this bond involves the back donation of electrons, the \({\rm{M}} \to {\rm{CO\pi }}\) bond is also known as the back bond.

The d-orbital of the metal atom involved in the \({\rm{M}} \to {\rm{CO\pi }}\) back bonding should have a proper alignment with the \(\left( {{\rm{\pi 2P}}{{\rm{x}}^{\rm{}}}{\rm{ = \pi 2P}}{{\rm{y}}^{\rm{}}}} \right)\) antibonding orbitals of the \({\rm{CO}}\) molecule. Hence, out of \({{\rm{d}}_{{\rm{xy}}}},{{\rm{d}}_{{\rm{yz}}}},{{\rm{d}}_{{\rm{zx}}}}\), orbitals, anyone orbital donates an electron to the empty \(\left( {{\rm{\pi 2P}}{{\rm{x}}^{\rm{}}}{\rm{ = \pi 2P}}{{\rm{y}}^{\rm{}}}} \right)\) antibonding orbitals of the \({\rm{CO}}\) molecule.

In metal carbonyls \(\left( {\sigma 2\;{{\rm{s}}^*}} \right)\) orbital of \({\rm{CO}}\) acts as a very weak donor, whereas \({{\rm{\pi }}^{\rm{*}}}\) orbitals act as electron acceptors. The lone pair of electrons on the carbon atom acts as the Lewis σ base (an electron-pair donor to form \(\sigma \) bond), and the empty \({\rm{CO}}\) antibonding orbital acts as the Lewis \({\rm{\pi }}\) acid (an electron pair acceptor to form \({\rm{\pi }}\) bond).

Carbon monoxide is one of the most important \({\rm{\pi }}\)- acceptor ligands. Because of its \({\rm{\pi }}\)- acidity, carbon monoxide can stabilize metals’ zero formal oxidation state in carbonyl complexes. In \({\rm{CO}}\),

1. The lower energy atomic orbitals of oxygen contribute more to the bonding molecular orbital.
2. The higher energy atomic orbital of carbon contributes more to the antibonding molecular orbitals.

As the HOMO \(\left( {\sigma 2\;{{\rm{s}}^*}} \right)\) and LUMO \(\left( {{\rm{\pi 2P}}{{\rm{x}}^{\rm{}}}{\rm{ = \pi 2P}}{{\rm{y}}^{\rm{}}}} \right)\) of the \({\rm{CO}}\) molecule involved in metal carbonyl bonding are the antibonding orbitals, primarily derived from a \({\rm{2 p}}\) orbital of the carbon atom. Hence, the lone pair of electrons in HOMO resides on the \({\rm{C}}\) atom.

Synergic Effect

The metal-carbon bond in metal carbonyls possesses both \({\rm{s}}\) and \({\rm{p}}\) character.

1. The overlapping of the empty hybrid orbital of the metal atom with the filled hybrid orbital (HOMO) of the carbon atom of carbon monoxide molecule results in the formation of an \({\rm{M}} \leftarrow {\rm{CO}}\sigma \)-bond.

2. And the overlapping of filled dπ orbitals or hybrid dpπ orbitals of the metal atom with low-lying empty (LUMO) orbitals of \({\rm{CO}}\) molecule results in the formation of \({\rm{M}} \to {\mathop{\rm CO}\nolimits} {\rm{\pi }}\) bonding. The metal to ligand bonding creates a synergic effect which strengthens the bond between \({\rm{CO}}\) and the central metal atom.

Evidence of Synergistic Bonding

The energy required to stretch a bond along the bond axis is measured by the Stretching frequencies. The higher the stretching frequency, the higher is the energy needed to stretch a bond along the bond axis.

Assumption: If the synergistic bonding model is valid, then we would expect the length and strength of the \({\rm{CO}}\) bond to be affected as electrons are pushed into the \({{\rm{\pi }}^{\rm{*}}}\) orbital. As the \({\rm{CO}}\) bond becomes weaker and longer, it should vibrate at a lower frequency.

Observation: The stretching frequency of \({\rm{CO}}\) molecule consisting of triple bond is around
\(2143\;{\rm{c}}{{\rm{m}}^{ – 1}}\) in the gaseous state. The stretching frequency of some of the metal carbonyls are found to be:

Metal Carbonyl	Stretching Frequency of \({\rm{CO}}\) bond
\({\left[ {{\rm{Fe}}{{({\rm{CO}})}_4}} \right]^{2 – }}\)	\(1790\;{\rm{c}}{{\rm{m}}^{ – 1}}\)
\({\left[ {{\rm{Co}}{{({\rm{CO}})}_4}} \right]^ – }\)	\(1890\;{\rm{c}}{{\rm{m}}^{ – 1}}\)
\({\rm{Ni}}{({\rm{CO}})_4}\)	\(2060\;{\rm{c}}{{\rm{m}}^{ – 1}}\)

The stretching frequencies of the above carbonyl complexes are found to be lesser than the stretching frequency of the \({\rm{CO}}\) molecule. Hence we can conclude that a synergistic bond exists in these metal complexes.

Consequences of Synergistic Bonding

1. Let us take the examples of the following metal carbonyls.

Metal Carbonyl	Stretching Frequency of \({\rm{CO}}\) bond
\({\left[ {{\rm{Fe}}{{({\rm{CO}})}_4}} \right]^{2 – }}\)	\(1790\;{\rm{c}}{{\rm{m}}^{ – 1}}\)
\({\left[ {{\rm{Co}}{{({\rm{CO}})}_4}} \right]^ – }\)	\(1890\;{\rm{c}}{{\rm{m}}^{ – 1}}\)
\({\rm{Ni}}{({\rm{CO}})_4}\)	\(2060\;{\rm{c}}{{\rm{m}}^{ – 1}}\)

In the above examples, the \({\left[ {{\rm{Fe}}{{({\rm{CO}})}_4}} \right]^{2 – }}\) carbonyl complex has the least stretching frequency and \({\rm{Ni}}{({\rm{CO}})_4}\) has the highest stretching frequency. The difference in the stretching frequencies of these metal complexes is due to the weakening of the \({\rm{CO}}\) bond and the strengthening of the \({\rm{ M – C}}\) bond.

Hence, \({\left[ {{\rm{Fe}}{{({\rm{CO}})}_4}} \right]^{2 – }}\) has the weakest \({\rm{CO}}\) bond and strongest \({\rm{MC}}\) bond, whereas \({\rm{Ni}}{({\rm{CO}})_4}\) has the strongest \({\rm{CO}}\) bond and weakest \({\rm{MC}}\) bond.

2. Moreover, the stretching frequency of the \({\rm{CO}}\) bond is closely related to the overall charge of the metal carbonyls. As \({\rm{CO}}\) is a neutral ligand, the overall charge of the metal carbonyl belongs to the central metal atom.

Metal carbonyl	Charge on the central metal atom
\({\left[ {{\rm{Fe}}{{({\rm{CO}})}_4}} \right]^{2 – }}\)	\(-2\)
\({\left[ {{\rm{Co}}{{({\rm{CO}})}_4}} \right]^ – }\)	\(-1\)
\({\rm{Ni}}{({\rm{CO}})_4}\)	\(0\)

Positive Oxidation State of the Central Metal Atom

The carbonyl ligand bonds with only those metals that have zero or less than \(+2\) oxidation state. The presence of a positive oxidation state indicates that the metal atom is deficient in electrons and back donating electrons from an electron-deficient species is difficult.

Hence, the lower the central metal atom’s oxidation state, the more favourable the back donation of electrons to the \({{\rm{\pi }}^{\rm{*}}}\) orbitals of the \({\rm{CO}}\) molecule. The overall charge of the metal carbonyls determines the extent to which a metal atom will back donate electrons and strengthen the \({\rm{ M – C}}\) bond.

Negative Oxidation State of the Central Metal Atom

A negative charge on the metal carbonyl denotes that the metal atom has an excess of electrons and can easily donate the electrons to \({\rm{CO}}\) for back bonding. This results in the strengthening of the \({\rm{ M – C}}\) bond and the weakening of the \({\rm{ C – O}}\) bond. The higher the negative charge on the metal carbonyl, the easier the back donation, and the stronger the \({\rm{ M – C}}\) bond.

Based on these facts, we can conclude that-

1. \({\left[ {{\rm{Fe}}{{({\rm{CO}})}_4}} \right]^{2 – }}\), being an electron-rich species can easily back donate electrons to the \({\rm{CO}}\) ligand. Among the three metal carbonyls, the \({\rm{Fe}} – {\rm{C}}\) bond is the strongest, and the \({\rm{ C – O}}\) bond is the weakest.

2. \({\left[ {{\rm{Co}}{{({\rm{CO}})}_4}} \right]^ – }\) is a uni negative complex and can also back donate electrons but not as easily as of the \({\left[ {{\rm{Fe}}{{({\rm{CO}})}_4}} \right]^{2 – }}\) metal carbonyl. The \({\rm{ Co – C }}\) bond is strong, but it is less than the \({\rm{Fe}} – {\rm{C}}\) bond.

3. In \({\rm{Ni}}{({\rm{CO}})_4}\) has zero overall charge, and hence is an electron-deficient species. Inspite of being an electron-deficient species, it donates electrons to back the bond with the carbon atom of the \({\rm{CO}}\) ligand. \({\rm{ Ni – C}}\) bond is the weakest, and the \({\rm{ C – O}}\) bond is the strongest among the three. This is quite evident from the \({\rm{CO}}\) stretching frequencies.

Thus, we can conclude that the lower the stretching frequency, the stronger and shorter the \({\rm{ M – C}}\) bond, the weaker and longer the \({\rm{ C – O}}\) bond.

The order of \({\rm{CO}}\) bond strengths of some metal carbonyl complexes is-

1. \({\left[ {{\rm{M(CO)6}}} \right]^{\rm{ + }}}{\rm{ > }}\left[ {{\rm{Cr(CO)6}}} \right]{\rm{ > }}{\left[ {{\rm{V(CO)6}}} \right]^{\rm{ – }}}{\rm{ > }}{\left[ {{\rm{Ti(CO)6}}} \right]^{{\rm{2 – }}}}\)

2. \({\rm{Ni(CO)4 > }}{\left[ {{\rm{Co(CO)4}}} \right]^{\rm{ – }}}{\rm{ > }}{\left[ {{\rm{Fe(CO}}{{\rm{)}}_{\rm{4}}}} \right]^{{\rm{2 – }}}}\)

Summary

The hallmark ligand of organometallic chemistry is the \({\rm{CO}}\) ligand. The \({\rm{CO}}\) ligands bind tightly to the metal centre using a synergistic mechanism that involves donating the \({\rm{CO}}\) ligand lone pair to the empty orbitals of metal. It is then followed by the \({\rm{\pi }}\)−back donation from a filled metal \({\rm{d}}\) orbital to a vacant \({\sigma ^*}\) orbital of the \({\rm{CO}}\) ligand. In this article, we learned the \({\rm{MO}}\) diagram of the \({\rm{CO}}\) molecule, which forms the base of bonding in metal carbonyls. We also learned the consequences of synergic bonding.

FAQs

Q.1. What is metal carbonyl bonding?
Ans: The bonding between the transition metal and carbonyl ligand in a coordination complex is called metal carbonyl bonding. It possesses both \({\rm{s}}\) and \({\rm{p}}\) character. This is because, in metal carbonyls, a sigma bond is formed by the donation of electrons from the \(\sigma 2\;{{\rm{s}}^*}\) HOMO orbital of the \({\rm{CO}}\) ligand to the vacant \({{\rm{d}}_{{{\rm{x}}^2} – {{\rm{y}}^2}}}\) orbital of the metal atom. A \({\rm{\pi }}\) bond is formed by the back donation of electrons from the \({{\rm{d}}_{{\rm{xy}}}}\) or \({{\rm{d}}_{{\rm{yz}}}}\) or \({{\rm{d}}_{{\rm{zx}}}}\) orbital of the metal atom present in the lowest possible oxidation state to the \(\left( {{\rm{\pi 2P}}{{\rm{x}}^{\rm{}}}{\rm{ = \pi 2P}}{{\rm{y}}^{\rm{}}}} \right)\) LUMO of the \({\rm{CO}}\) ligand. This results in back bonding and is known as synergistic bonding.

Q.2. Which carbonyl has the strongest \({\rm{CO}}\) Bond?
Ans: The metal carbonyls having the weakest \({\rm{M-C}}\) bond have the strongest \({\rm{CO}}\) bond.

Q.3. What is the synergic effect in metal carbonyls?
Ans: Metal carbonyls have both \({\rm{s}}\) and \({\rm{p}}\) character in their metal-carbon bonds. The overlapping of empty hybrid orbital on metal atom with the filled hybrid orbital (HOMO) on carbon atom of carbon monoxide molecule results into the formation of a \({\rm{M}} \leftarrow {\rm{CO}}\sigma \)-bond. And the overlapping of filled \({\rm{d\pi }}\) orbitals or hybrid \({\rm{dp\pi }}\) orbitals of the metal atom with low-lying empty (LUMO) orbitals on \({\rm{CO}}\) molecule results in the formation of \({\rm{M}} \to {\rm{CO\pi }}\) bonding. The synergic impact of metal to ligand bonding increases the link between \({\rm{CO}}\) and the metal.

Q.4. What are the different types of bonds present in metal carbonyls?
Ans: The metal-carbon bond in metal carbonyls has characteristics of both \(\sigma \) and \({\rm{\pi }}\) bonds. The σ bond is formed by the donation of electrons from the HOMO of the \({\rm{CO}}\) molecule to the empty orbitals of the metal atom whereas the \({\rm{\pi }}\) bond is formed by the back donation of electrons from filled d-orbitals to the LUMO of the \({\rm{CO}}\) molecule. The synergic effect strengthens the \({\rm{M-C}}\) bond in metal carbonyls.

Q.5. What factors affect the back-bonding phenomenon?
Ans: Factors that affect the back bonding phenomenon includes:
1. Charge on the metal- A negative charge strengthens the synergistic bond, whereas a high positive charge makes back donation difficult.
2. Contribution of other ligands on the metal centre-electron donating ligands increase the back-bonding process while the electron-withdrawing ligands inhibit the back-bonding process.