In \(1884,\) Ludwig Mond observed that the nickel valves were being eaten away by \({\rm{CO}}\) gas in a nickel refining industry. He heated nickel powder in a stream of \({\rm{CO}}\) gas to synthesize the first known metal carbonyl compound in the form of \({\rm{Ni}}{\left( {{\rm{CO}}} \right)_{\rm{4}}}{\rm{.}}\)

The Mond refining process was thus born in which he proposed that the volatile \({\rm{Ni}}{\left( {{\rm{CO}}} \right)_{\rm{4}}}\) compound can be decomposed to pure metal at elevated temperature. In this article, we will provide detailed information on Metal Carbonyls. This article will also discuss uses of metal carbonyls, transition metal carbonyls, synergic bond in metal carbonyls, bonding in metal carbonyls and example of metal carbonyls.

What are Metal Carbonyls?

Metal carbonyls are coordination complexes of transition metals with carbon monoxide ligands. The general formula of metal carbonyls is \({{\rm{M}}_{\rm{x}}}{\left( {{\rm{CO}}} \right)_{\rm{y}}}{\rm{.}}\) A lone pair of electrons are available on both carbon and oxygen atoms of a carbon monoxide ligand. As the carbon atoms donate electrons to the metal, these complexes are named carbonyls.

Classification of Metal Carbonyls

1.Based on the type of Ligands.
Metal carbonyls can be classified into the following, based on the type of ligand they are attached to. These are-
a. Homoleptic Metal carbonyls: In homoleptic metal carbonyls, only one type of ligand is present. These metal carbonyls have only \({\rm{CO}}\) ligands with the general formula \({{\rm{M}}_{\rm{x}}}{\left( {{\rm{CO}}} \right)_{\rm{y}}}{\rm{.}}\)
Example – \({\rm{Ni}}{\left( {{\rm{CO}}} \right)_{\rm{4}}}\)
b. Heteroleptic Metal carbonyls: In heteroleptic metal carbonyls, more than one type of ligands are present. Along with \({\rm{CO}}\) ligands, other types of ligands are also attached to the metal. These have the general formula \({{\rm{M}}_{\rm{x}}}{{\rm{L}}_{\rm{m}}}{\left( {{\rm{CO}}} \right)_{\rm{y}}}{\rm{.}}\)
Example – \(\left[ {{\rm{Pt}}{{\left( {{\rm{CO}}} \right)}_{\rm{2}}}{\rm{C}}{{\rm{l}}_{\rm{2}}}} \right]\)

2. Based on the number of metal centres.

Metal carbonyls can be classified into the following based on the number of metal centres. These are-
a. Mononuclear metal carbonyl: These metal carbonyls have only one metal centre. The general formula of mononuclear metal carbonyls is \({\rm{M}}{\left( {{\rm{CO}}} \right)_{\rm{n}}}{\rm{.}}\) Example – \({\rm{Ni}}{\left( {{\rm{CO}}} \right)_{\rm{4}}}\)
b. Polynuclear metal carbonyl – These metal carbonyls have two or more metal centres. The general formula of polynuclear metal carbonyls is \({{\rm{M}}_{\rm{x}}}{{\rm{N}}_{\rm{y}}}{\left( {{\rm{CO}}} \right)_{\rm{n}}}{\rm{.}}\) Example – \({\rm{M}}{{\rm{n}}_{\rm{2}}}{\left( {{\rm{CO}}} \right)_{{\rm{10}}}}{\rm{,}}\,{\rm{MnCo}}{\left( {{\rm{CO}}} \right)_{\rm{9}}}\)

Based on the type of metals, the polynuclear metal carbonyls can be classified into-
a. Homonuclear: These metal carbonyls have two or more than two metal centres but of the same type. Example – \({\rm{M}}{{\rm{n}}_{\rm{2}}}{\left( {{\rm{CO}}} \right)_{{\rm{10}}}}\)
b. Heteronuclear: These metal carbonyls have two or more than two metal centres but of different types. Example- \({\rm{MnCo}}{\left( {{\rm{CO}}} \right)_{\rm{9}}}\)
Based on bridge ligands, the polynuclear metal carbonyls can be classified into-
a. \({{\rm{\mu }}_{\rm{2}}}\) – bridging: In this type of metal carbonyl, the \({\rm{CO}}\) ligand is connected to two metal centres. Example- \({\rm{F}}{{\rm{e}}_{\rm{2}}}{\left( {{\rm{CO}}} \right)_{\rm{9}}}\)
b. \({{\rm{\mu }}_3}\) – bridging: In this type of metal carbonyl, the CO ligand is connected to three metal centres. Example- \(\left( {{\rm{c}}{{\rm{p}}_{\rm{4}}}} \right){\rm{F}}{{\rm{e}}_{\rm{4}}}{\left( {{\rm{CO}}} \right)_{\rm{4}}}\)
c. Non-bridging: In this type of metal carbonyls, there are no bridged \({\rm{CO}}\) ligands. Only terminal \({\rm{CO}}\) ligands are present. Example : \({\rm{M}}{{\rm{n}}_{\rm{2}}}{\left( {{\rm{CO}}} \right)_{{\rm{10}}}}\)

Preparation of Metal Carbonyl Complexes

Metal carbonyls can be prepared in the following ways:

I. Directly Using the \(^{\rm{‘}}{\rm{C}}{{\rm{O}}^{\rm{‘}}}{\rm{.}}\)

Mononuclear metal carbonyls such as \({\rm{Fe}}{\left( {{\rm{CO}}} \right)_{\rm{2}}}{\rm{,}}\,{\rm{Ni}}{\left( {{\rm{CO}}} \right)_{\rm{4}}}\) and binuclear metal carbonyls such as \({\rm{C}}{{\rm{o}}_{\rm{2}}}{\left( {{\rm{CO}}} \right)_{\rm{8}}}\) maybe prepared by direct reaction of \({\rm{CO}}\) with finely divided metal at suitable temperature and pressure.

The main requirement of this method is that the metal centre must be in a reduced low oxidation state to facilitate \({\rm{CO}}\) binding to the metal centre through metal to ligand \(\pi – \)back donation.

II. Reductive Carbonylation

Treating salts like \({\rm{Ru}}{\left( {{\rm{acac}}} \right)_{\rm{3}}}{\rm{,}}\,{\rm{CrC}}{{\rm{l}}_{\rm{3}}}{\rm{,}}\,{\rm{VC}}{{\rm{l}}_{\rm{3}}}{\rm{,}}\,{\rm{CoS,}}\,{\rm{Co}}{{\rm{l}}_2}\) with carbon monoxide in the presence of suitable reducing agents like \({\rm{Mg,}}\,{\rm{Ag,}}\,{\rm{Cu,}}\,{\rm{Na,}}\,{{\rm{H}}_{\rm{2}}}{\rm{,}}\,{\rm{LiAl}}{{\rm{H}}_{\rm{4}}}{\rm{,}}\) etc. gives metal carbonyls.

III. From Mononuclear Carbonyls

When a cold solution of \({\rm{Fe}}{\left( {{\rm{CO}}} \right)_{\rm{5}}}{\rm{/Os}}{\left( {{\rm{CO}}} \right)_{\rm{5}}}\) in glacial acetic acid \({\rm{C}}{{\rm{H}}_{\rm{3}}}{\rm{COOH}}\) is irradiated with ultraviolet light, \({\rm{F}}{{\rm{e}}_{\rm{2}}}{\left( {{\rm{CO}}} \right)_{\rm{9}}}{\rm{/O}}{{\rm{s}}_{\rm{2}}}{\left( {{\rm{CO}}} \right)_{\rm{9}}}\) are obtained.

Structure of Metal Carbonyl

To understand the bonding in metal carbonyls, let us first see the \({\rm{MO}}\) diagram of carbon monoxide. The order of energy of the molecular orbitals and the accommodation of fourteen electrons of the carbon monoxide can be shown as:

The molecular orbital configuration of the \({\rm{CO}}\) molecule is:
\({\left( {{\rm{\sigma 1}}{{\rm{s}}^{\rm{b}}}} \right)^{\rm{2}}}{\left( {{\rm{\sigma 1}}{{\rm{s}}^{\rm{*}}}} \right)^{\rm{2}}}{\left( {{\rm{\sigma 2}}{{\rm{s}}^{\rm{b}}}} \right)^{\rm{2}}}{\left( {{\rm{\pi 2P}}{{\rm{x}}^{\rm{b}}}{\rm{ = \pi 2P}}{{\rm{y}}^{\rm{b}}}} \right)^{\rm{4}}}{\left( {{\rm{\sigma 2P}}{{\rm{z}}^{\rm{b}}}} \right)^{\rm{2}}}{\left( {{\rm{\sigma 2}}{{\rm{s}}^{\rm{*}}}} \right)^{\rm{2}}}\left( {{\rm{\pi 2P}}{{\rm{x}}^{\rm{*}}}{\rm{ = }}} \right.{\left. {{\rm{\pi 2P}}{{\rm{y}}^{\rm{*}}}} \right)^{\rm{0}}}{\left( {{\rm{\sigma 2P}}{{\rm{z}}^{\rm{*}}}} \right)^{\rm{0}}}\)
\({{\rm{\sigma 2}}{{\rm{s}}^{\rm{*}}}}\) is the highest occupied molecular orbital \(\left( {{\rm{HOMO}}} \right)\) that can donate the lone pair of electrons to form an \(\left( {{\rm{OC}} \to {\rm{M}}} \right)\sigma \) bond. \(\left( {{\rm{\pi 2P}}{{\rm{x}}^{\rm{*}}}{\rm{ = \pi 2P}}{{\rm{y}}^{\rm{*}}}} \right)\) are the lowest unoccupied molecular orbitals \(\left( {{\rm{LUMO}}} \right)\) that can accept the electron density from an appropriately oriented, filled metal orbital resulting in the formation of an \(\left( {{\rm{M}} \to {\rm{CO}}} \right){\rm{\pi }}\) bond.

The nature of \({\rm{M}} – {\rm{CO}}\) bonding in mononuclear carbonyls can be understood by considering the formation of a dative \(\sigma – \)bond and \({\rm{\pi }} – \)bond due to back donation.

Formation of Dative \(\sigma  – \)bond

The overlapping of empty hybrid orbital on a metal atom with the filled hybrid orbital \(\left( {{\rm{HOMO}}} \right)\) on the carbon atom of carbon monoxide molecule results into the formation of a \(\left( {{\rm{M}} \leftarrow {\rm{CO}}} \right)\sigma – \)bond. The central metal in metal carbonyls is a transition element. Hence, there is a vacant \({\rm{d}} – \)orbital available on the metal atom to accept the lone pair of electrons from the \({\rm{HUMO}}\) of the \({\rm{CO}}\) molecule.

As we know, there are five \({\rm{d}} – \)orbitals, \({{\rm{d}}_{{\rm{xy}}}}{\rm{,}}\,{{\rm{d}}_{{\rm{yz}}}}{\rm{,}}\,{{\rm{d}}_{{\rm{zx}}}}{\rm{,}}\,{{\rm{d}}_{{{\rm{y}}^{\rm{2}}}}}{\rm{,}}\,{{\rm{d}}_{{{\rm{x}}^{\rm{2}}}{\rm{ – }}{{\rm{y}}^{\rm{2}}}}}{\rm{,}}\) available on the metal atom that can participate in the \(\left( {{\rm{M}} \leftarrow {\rm{CO}}} \right){\rm{\sigma }} – \)bond, the orbital with the proper alignment is involved in the \(\left( {{\rm{M}} \leftarrow {\rm{CO}}} \right){\rm{\sigma }} – \)bond.

The donation of electrons takes place from the sigma orbital \(\left( {{\rm{\sigma 2}}{{\rm{s}}^{\rm{*}}}} \right)\) of the \({\rm{CO}}\) molecule. Hence, the \({\rm{d}} – \)orbital having an alignment similar to the sigma orbital participates in the bonding. From the image above, we can figure out that \({{\rm{d}}_{{{\rm{x}}^2} – {{\rm{y}}^2}}}\) Orbital has a similar alignment as that of the \({\rm{s}}\) orbital.

The \({{\rm{d}}_{{{\rm{x}}^2} – {{\rm{y}}^2}}}\) orbital is along the \({\rm{XY}}\) axis, which is in alignment with the \({\rm{s}}\) orbital. Hence, out of the five d orbitals, \({{\rm{d}}_{{{\rm{x}}^2} – {{\rm{y}}^2}}}\) participates in the \(\left( {{\rm{M}} \leftarrow {\rm{CO}}} \right)\,\sigma – \) bonding. The \(\left( {{\rm{M}} \leftarrow {\rm{CO}}} \right)\,\sigma – \) bond is diagrammatically represented as below-

Formation of \({\rm{\pi }} – \)bond by Back Donation

This bond is formed because of overlapping of filled \({\rm{d\pi }}\) orbitals or hybrid \({\rm{dp\pi }}\) orbitals of a metal atom with low-lying empty \(\left( {{\rm{LUMO}}} \right)\) orbitals on \({\rm{CO}}\) molecule. i.e. \(\left( {{\rm{M}} \to {\rm{CO}}} \right){\rm{\pi }}\) bonding. \(\left( {{\rm{\pi 2P}}{{\rm{x}}^*} = {\rm{\pi 2P}}{{\rm{y}}^*}} \right)\) are the lowest unoccupied molecular orbitals \(\left( {{\rm{LUMO}}} \right)\) that can accept the electron density from an appropriately oriented, filled metal orbital resulting in the formation of an \(\left( {{\rm{M}} \to {\rm{CO}}} \right){\rm{\pi }}\) bond.

This type of bonding takes place through a back donation of electrons, also known as back bonding. The \({\rm{d}} – \)orbital of the metal atom involved in the \(\left( {{\rm{M}} \to {\rm{CO}}} \right){\rm{\pi }}\) back bonding should have a proper alignment with the \(\left( {{\rm{\pi 2P}}{{\rm{x}}^*} = {\rm{\pi 2P}}{{\rm{y}}^*}} \right)\) antibonding orbitals of the \({\rm{CO}}\) molecule. Hence \({{\rm{d}}_{{\rm{xy}}}}\) orbital of the metal atom donates an electron to the empty \(\left( {{\rm{\pi 2P}}{{\rm{x}}^*} = {\rm{\pi 2P}}{{\rm{y}}^*}} \right)\) antibonding orbitals of the \({\rm{CO}}\) molecule.

In metal carbonyls \(\left( {{\rm{\sigma 2}}{{\rm{s}}^{\rm{*}}}} \right)\) orbital of \({\rm{CO}}\) acts as a very weak donor, and \({{\rm{\pi }}^{\rm{*}}}\) orbitals act as acceptors. The lone pair on the carbon atom acts as the Lewis \({\rm{\sigma }}\) base (an electron-pair donor to form \({\rm{\sigma }}\) 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 ligand. Because of its \({\rm{\pi – }}\) acidity, carbon monoxide can stabilize the zero oxidation state of metals in carbonyl complexes. In \({\rm{CO,}}\) the lower energy atomic orbitals of oxygen contribute more to the bonding molecular orbital. The higher energy atomic orbital of carbon contributes more to the antibonding molecular orbitals. As the \({\rm{HUMO}}\) and \({\rm{LUMO}}\) in \({\rm{CO}}\) molecule involved in metal carbonyl bonding are the antibonding orbitals, they are primarily derived from a \({\rm{2p}}\) orbital of the carbon atom. Hence, the lone pair of electrons in \({\rm{HUMO}}\) resides on the \({\rm{C}}\) atom.

Synergic Effect

The metal-carbon bond in metal carbonyls possesses both \({\rm{s}}\) and \({\rm{p}}\) character. The overlapping of empty hybrid orbital on a metal atom with the filled hybrid orbital \(\left( {{\rm{HOMO}}} \right)\) on the carbon atom of carbon monoxide molecule results in the formation of an \(\left( {{\rm{M}} \leftarrow {\rm{CO}}} \right)\,{\rm{\sigma }} – \) bond.

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

Evidence of Synergic Bonding

Stretching frequency measures the energy required to stretch a bond along the bond axis. The higher the energy needed to stretch a bond along the bond axis, the higher is the stretching frequency.
Bond Strength is directly proportional to the Stretching Frequency.

Assumption: If the synergic 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 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 \({\rm{2143}}\,{\rm{c}}{{\rm{m}}^{{\rm{ – 1}}}}\) 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
\({\left[ {{\rm{Fe}}{{\left( {{\rm{CO}}} \right)}_{\rm{4}}}} \right]^{{\rm{2 – }}}}\)	\(1790\;{\rm{c}}{{\rm{m}}^{ – 1}}\)
\({\left[ {{\rm{Co}}{{\left( {{\rm{CO}}} \right)}_{\rm{4}}}} \right]^ – }\)	\(1890\;{\rm{c}}{{\rm{m}}^{ – 1}}\)
\[{\rm{Ni}}{\left( {{\rm{CO}}} \right)_{\rm{4}}}\]	\(2060\;{\rm{c}}{{\rm{m}}^{ – 1}}\)

All the above metal carbonyls exhibit a stretching frequency lesser than that of the \({\rm{CO}}\) molecule. Hence we can conclude that a synergic bond exists.

Consequences of synergic bonding

Considering the above examples, the \({\left[ {{\rm{Fe}}{{\left( {{\rm{CO}}} \right)}_{\rm{4}}}} \right]^{{\rm{2 – }}}}\) metal carbonyl has the least stretching frequency and \({\rm{Ni}}{\left( {{\rm{CO}}} \right)_{\rm{4}}}\) has the highest stretching frequency. The difference exists due to the weakening of the \({\rm{CO}}\) bond and the strengthening of the \(\left( {{\rm{M}} – {\rm{C}}} \right)\) bond.

From the table above, we can observe that the stretching frequency of the \({\rm{CO}}\) bond is closely related to the overall charge of the metal carbonyl. \({\rm{CO}}\) is a neutral ligand, so the overall charge belongs to the central metal atom.

Metal Carbonyl	Charge on the central metal atom
\({\left[ {{\rm{Fe}}{{\left( {{\rm{CO}}} \right)}_{\rm{4}}}} \right]^{2 – }}\)	\( – 2\)
\({\left[ {{\rm{Co}}{{\left( {{\rm{CO}}} \right)}_{\rm{4}}}} \right]^{2 – }}\)	\( – 1\)
\({\rm{Ni}}{\left( {{\rm{CO}}} \right)_{\rm{4}}}\)	\(0\)

We know that the carbonyl ligand, bonds with a metal having zero or less than \( + 2\) oxidation state. This is because the metal has to back donate electrons to the \({{\rm{\pi }}^{\rm{*}}}\) orbitals of the \({\rm{CO}}\) molecule. The presence of a positive oxidation state denotes that the metal atom is deficient in electrons. Back donating electrons from an electron-deficient species is difficult. The extent to which a metal atom will back donate electrons and strengthen the \(\left( {{\rm{M}} – {\rm{C}}} \right)\) bond is measured by the overall charge of the metal carbonyls.

A negative charge on the metal carbonyl complex 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 \(\left( {{\rm{M}} – {\rm{C}}} \right)\) bond and the weakening of the \({\rm{C}} – {\rm{O}}\) bond.
The higher the negative charge on the metal carbonyl, the easier the back donation, and the stronger the \(\left( {{\rm{M}} – {\rm{C}}} \right)\) bond.

Based on this fact, \({\left[ {{\rm{Fe}}{{\left( {{\rm{CO}}} \right)}_{\rm{4}}}} \right]^{2 – }},\) being electron-rich species can easily back donate electrons to the \({\rm{CO}}\) ligand. The \({\rm{Fe}} – {\rm{C}}\) bond is the strongest, and the \({\rm{C}} – {\rm{O}}\) bond is the weakest among the three metal carbonyls. \({\left[ {{\rm{Co}}{{\left( {{\rm{CO}}} \right)}_4}} \right]^ – }\) has only one negative charge and can also back donate electrons but not as easily as of the \({\left[ {{\rm{Fe}}{{\left( {{\rm{CO}}} \right)}_{\rm{4}}}} \right]^{2 – }}\) metal carbonyl. The \({\rm{Co}} – {\rm{C}}\) bond is strong, but it is less than the \({\rm{Fe}} – {\rm{C}}\) bond.

In \({\rm{Ni}}{\left( {{\rm{CO}}} \right)_{\rm{4}}}\) as the metal carbonyl has zero charges, it is an electron-deficient species. In spite of being an electron-deficient species, it donates electrons to form back bonding with the carbon atom of the \({\rm{CO}}\) ligand. \(\left( {{\rm{Ni}} – {\rm{C}}} \right)\) bond is the weakest, and the \({\rm{C}} – {\rm{O}}\) 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 \({{\rm{M}} – {\rm{C}}}\) bond, the weaker and longer the \({\rm{C}} – {\rm{O}}\) bond.

The order of \({\rm{CO}}\) bond strengths is-
(i) \({\left[ {{\rm{M}}{{\left( {{\rm{CO}}} \right)}_{\rm{6}}}} \right]^{\rm{ + }}}{\rm{ > }}\left[ {{\rm{Cr}}{{\left( {{\rm{CO}}} \right)}_{\rm{6}}}} \right]{\rm{ > }}{\left[ {{\rm{V}}{{\left( {{\rm{CO}}} \right)}_{\rm{6}}}} \right]^{\rm{ – }}}{\rm{ > }}{\left[ {{\rm{Ti(CO}}{{\rm{)}}_{\rm{6}}}} \right]^{{\rm{2 – }}}}\)
(ii) \({\rm{Ni}}{\left( {{\rm{CO}}} \right)_{\rm{4}}}{\rm{ > }}{\left[ {{\rm{Co}}{{\left( {{\rm{CO}}} \right)}_{\rm{4}}}} \right]^ – } > {\left[ {{\rm{Fe}}{{\left( {{\rm{CO}}} \right)}_{\rm{4}}}} \right]^{2 – }}\)

Examples of Metal carbonyls in coordination compounds
1. \({\rm{Ni}}{\left( {{\rm{CO}}} \right)_{\rm{4}}},\) Nickel Tetracarbonyl
In \({\rm{Ni}}{\left( {{\rm{CO}}} \right)_{\rm{4}}},\) Nickel atom is \({\rm{s}}{{\rm{p}}^{\rm{3}}}\) hybridized. Ten electrons are present in the valence shell of the \({\rm{Ni}}\) atom \(\left( {{\rm{Ni}} = {\rm{3}}{{\rm{d}}^{\rm{8}}}{\rm{4}}{{\rm{s}}^{\rm{2}}}} \right)\) that get paired in \({\rm{3d}}\) orbitals. Thus, the valence shell configuration of the \({\rm{Ni}}\) atom in \({\rm{Ni}}{\left( {{\rm{CO}}} \right)_{\rm{4}}},\) molecule becomes \({\rm{3}}{{\rm{d}}^{{\rm{10}}}}{\rm{4}}{{\rm{s}}^{\rm{0}}}{\rm{.}}\) The absence of unpaired electrons makes \({\rm{Ni}}{\left( {{\rm{CO}}} \right)_{\rm{4}}},\) diamagnetic.

The \(\left( {{\rm{OC}} \to {\rm{Ni}}} \right)\) bond results from the overlap between the empty \({\rm{s}}{{\rm{p}}^{\rm{3}}}\) hybrid orbitals on the \({\rm{Ni}}\) atom and the \({\rm{HOMO}}\) on the \({\rm{C}}\) atom in the \({\rm{CO}}\) molecule

Structure: Nickel tetracarbonyl has a tetrahedral geometry with \({\rm{Ni}} – {\rm{C}}\) bond lengths of \(1.5\,\) Å.

2. \({\rm{Fe}}{\left( {{\rm{CO}}} \right)_{\rm{5}}}{\rm{,}}\) Iron pentacarbonyl In \({\rm{Fe}}{\left( {{\rm{CO}}} \right)_{\rm{5}}}{\rm{Fe}}\) atom is \({\rm{ds}}{{\rm{p}}^{\rm{3}}}\) hybridized. Eight electrons present in the valence shell of the \({\rm{Fe}}\) atom \(\left( {{\rm{Fe}}:{\rm{3}}{{\rm{d}}_6}{\rm{4}}{{\rm{s}}_2}} \right)\) that get paired in four \({\rm{3d}}\) orbitals.

Thus the valence shell configuration of \({\rm{Fe}}\) in \({\rm{Fe}}{\left( {{\rm{CO}}} \right)_{\rm{5}}}\) becomes \({\rm{3}}{{\rm{d}}^{\rm{8}}}{\rm{4}}{{\rm{s}}^{\rm{0}}}{\rm{.}}\) The \(\left( {{\rm{OC}} \to {\rm{Fe}}} \right)\) bond results from the overlap between the empty \({\rm{ds}}{{\rm{p}}^{\rm{3}}}\) hybrid orbitals on the \({\rm{Fe}}\) atom and the \({\rm{HOMO}}\) on the \({\rm{C}}\) atom in the \({\rm{CO}}\) molecule.

Structure: Structural studies have suggested trigonal bipyramidal geometry for iron pentacarbonyl. The \({\rm{Fe}} – {\rm{C}}\) distances are \(1.84\) Å and \(1.80\) Å for axial and equatorial bonds, respectively. The absence of unpaired electrons makes \({\rm{Fe}}{\left( {{\rm{CO}}} \right)_{\rm{5}}}\) diamagnetic.

3. \({\rm{Cr}}{\left( {{\rm{CO}}} \right)_{\rm{6}}},\) Chromium hexacarbonyl:
In \({\rm{Cr}}{\left( {{\rm{CO}}} \right)_{\rm{6}}}{\rm{,}}\,{\rm{Cr}}\) atom is \({{\rm{d}}^{\rm{2}}}{\rm{s}}{{\rm{p}}^{\rm{3}}}\) hybridized. There are six electrons present in the valence shell of the \({\rm{Cr}}\) atom \(\left( {{\rm{Cr:3}}{{\rm{d}}^{\rm{5}}}{\rm{4}}{{\rm{s}}^{\rm{1}}}} \right)\) paired in three \({\rm{3d}}\) orbitals.

Thus the valence shell configuration of \({\rm{Cr}}\) in \({\rm{Cr}}{\left( {{\rm{CO}}} \right)_{\rm{6}}}\) becomes \(3\;{{\rm{d}}^6}{45^0}.\) The \(\left( {{\rm{OC}} \to {\rm{Cr}}} \right)\) bond results from the overlap between the empty \({{\rm{d}}^{\rm{2}}}{\rm{s}}{{\rm{p}}^{\rm{3}}}\) hybrid orbitals on the \({\rm{Cr}}\) atom and the \({\rm{HOMO}}\) on the \({\rm{C}}\) atom in the \({\rm{CO}}\) molecules.

Structure: Structural studies have suggested an octahedral geometry for chromium hexacarbonyl. The \(\left( {{\rm{Cr}} – {\rm{C}}} \right)\) bond length is found to be \(1.92\) Å while the \(\left( {{\rm{C}} – {\rm{O}}} \right)\) bond length is \(1.16\) Å The absence of unpaired electrons makes \({\rm{Cr}}{\left( {{\rm{CO}}} \right)_{\rm{6}}}\) diamagnetic.

4. \({\rm{M}}{{\rm{n}}_{\rm{2}}}{\left( {{\rm{CO}}} \right)_{{\rm{10}}}}{\rm{,}}\) Dimanganese decacarbonyl:
Manganese with an odd atomic number \(\left( {{\rm{Z}} = 25} \right)\) does not form manganese pentacarbonyl. Dimanganese decacarbonyl consists of two manganese pentacarbonyl groups joined through an \({\rm{Mn}} – {\rm{Mn}}\,\left( {2.79} \right) Å\) bond. The formation of this intermetallic bond effectively adds one electron to each of the manganese atoms. Thus, manganese, an element with an odd atomic number, forms a binuclear carbonyl compound. Since the molecule does not have any unpaired electrons, it is diamagnetic. The remaining two members of group \({\rm{VIIB}}\) viz. Technetium \(\left( {{\rm{Tc}}} \right)\) and Rhenium \(\left( {{\rm{Re}}} \right)\) also form decarbonyls with similar structures.

5. \({\rm{C}}{{\rm{o}}_{\rm{2}}}{\left( {{\rm{CO}}} \right)_{\rm{8}}}{\rm{,}}\) Dicobalt octacarbonyl:
Dicobalt octacarbonyl is known to exist in two isomeric forms. The two isomeric forms are-
d. A bridged structure of \({\rm{C}}{{\rm{o}}_{\rm{2}}}{\left( {{\rm{CO}}} \right)_{\rm{8}}}\) molecule is observed in the solid-state as well as solution state at a very low temperature. In the bridged structure, the cobalt atoms are in a \({{\rm{d}}^{\rm{2}}}{\rm{s}}{{\rm{p}}^{\rm{3}}}\) hybrid state.

Three such hybrid orbitals on each cobalt atom accept lone pair of electrons from three carbon monoxide molecules to form six \(\left( {{\rm{Co}} \to {\rm{CO}}} \right)\) coordinate bonds. A \({\rm{Co}} – {\rm{Co}}\) bond is formed by the overlapping of two half-filled \({{\rm{d}}^{\rm{2}}}{\rm{s}}{{\rm{p}}^{\rm{3}}}\) hybrid orbitals on the cobalt atoms. The remaining two half-filled hybrid orbitals on each \({\rm{Co}}\) atom overlap with appropriate orbital on the carbonyl carbon atom to form two bridging \({\rm{CO}}\) groups. Thus, all electrons in this molecule are paired, and it is diamagnetic.

e. A non – bridged structure of \({\rm{C}}{{\rm{o}}_{\rm{2}}}{\left( {{\rm{CO}}} \right)_{\rm{8}}}\) predominates in a solution at temperatures above ambience. In this structure, the cobalt atoms are in a \({\rm{ds}}{{\rm{p}}^{\rm{3}}}\) hybrid state. Out of the five hybrid orbitals on each cobalt atom, four orbitals on each cobalt atom accept a lone pair of electrons from the carbon monoxide molecules to form eight \(\left( {{\rm{Co}} – {\rm{CO}}} \right)\) coordinate bonds. One half-filled orbital on each cobalt atom overlaps to form a \({\rm{Co}} – {\rm{Co}}\) bond.

6. \({\rm{F}}{{\rm{e}}_{\rm{2}}}{\left( {{\rm{CO}}} \right)_{\rm{9}}}{\rm{,}}\) Diiron nonacarbonyl:
Each of the iron atoms in diiron nonacarbonyl has three terminal carbonyl groups. The remaining three carbon monoxide ligands act as \({\rm{\mu 2}} – {\rm{CO}}\) groups. In addition to this, there is a weak \({\rm{Fe}} – {\rm{Fe}}\) bond \(\left( {{\rm{2}}{\rm{.46}}} \right) Å\) formed by sharing of two unpaired electrons present in the \({\rm{3d}}\) orbitals of iron atoms.

Thus, both the iron atoms in the molecule are identical with coordination number seven. Since the molecule does not have any unpaired electron, it is diamagnetic. The structure of this molecule can be explained using \({{\rm{d}}^{\rm{2}}}{\rm{s}}{{\rm{p}}^{\rm{3}}}\) hybridization in \({\rm{Fe}}\) atoms, as shown in the figure.

7. \({\rm{F}}{{\rm{e}}_{\rm{3}}}{\left( {{\rm{CO}}} \right)_{{\rm{12}}}}{\rm{,}}\) Triiron dodecacarbonyl:
Triiron dodecacarbonyl has a \(3\)-membered ring structure. Two iron atoms in this molecule have three terminal carbonyl groups, while the third iron atom is connected to four terminal carbonyls. Two \({\rm{\mu 2}} – {\rm{CO}}\) groups also connect the former iron atoms. In addition to this, there are three \({\rm{Fe}} – {\rm{Fe}}\) bonds \[\left( {2.8} \right) Å\] connecting each of the iron atoms.

What are the Properties of Metal Carbonyls?

1. State: The majority of metallic carbonyls are liquids or volatile solids.
2. Colour: Most of the mononuclear carbonyls are colourless to pale yellow. \({\rm{V}}{\left( {{\rm{CO}}} \right)_6}\) is a bluish-black solid. Polynuclear carbonyls are dark in colour.
3. Solubility: Metal carbonyls are soluble in organic solvents like glacial acetic acid, acetone, benzene, carbon tetrachloride and ether.
4. Toxicity: Due to low melting points and poor thermal stability, they show toxicity related to the corresponding metal and carbon monoxide. Exposure to these compounds can cause damage to the lungs, liver, brain and kidneys. Nickel tetracarbonyl exhibits the strongest inhalation toxicity. These compounds are carcinogenic over long-term exposure.
5. Magnetic Property: The metals with even atomic number form mononuclear carbonyls. Metals with odd atomic number form dinuclear metal carbonyls. Thus, all the electrons in the metal atoms are paired. The unpaired electrons are utilized for the formation of metal-metal bonds. Hence these are generally diamagnetic in nature.
6. Thermal Stability: Most of the metal carbonyls melt or decompose at low temperatures. Solid carbonyls sublime in a vacuum, but they undergo some degree of degradation.
7. Thermodynamic Stability: Metal carbonyls are thermodynamically unstable. They undergo aerial oxidation at different rates. \({\rm{C}}{{\rm{o}}_2}{\left( {{\rm{CO}}} \right)_8}\) and \({\rm{F}}{{\rm{e}}_2}{\left( {{\rm{CO}}} \right)_9}\) are oxidized by air at room temperature, while chromium and molybdenum hexacarbonyl is oxidized in the air when heated.

Uses of Metal Carbonyls

The uses of metal carbonyls are listed below:

1. Metal carbonyls like carbonyl of iron are used for the preparation of inductors, pigments as dietary supplements.
2. Metal carbonyls are used as a catalyst in the production of a number of industrially important compounds, such as aldehydes.
3. The oxo-aldehydes resulting from hydroformylation are used for the large-scale synthesis of fatty alcohols, which are precursors to detergents.
4. Metal carbonyl complexes are being developed as potential drugs to release \({\rm{CO}}{\rm{.}}\) At low concentrations, \({\rm{CO}}\) functions as a vasodilator and an anti-inflammatory agent.

Summary

Metal carbonyls form an important class of coordination compounds. As these are mostly used as catalysts, their knowledge is vital to the production of many organic and inorganic compounds. Through this article, we have learned the various types of metal carbonyls and their related structure. We also learned the bonding between the metal carbonyls and their uniqueness from the rest of the coordination compounds.

NCERT Solutions for 12th Chemistry Chapter 9

Frequently asked questions (FAQs) on Metal Carbonyls

Q.1: What are metal carbonyls examples?
Ans: Coordination complexes of transition metals such as \(\left[ {{\rm{Fe}}{{\left( {{\rm{CO}}} \right)}_{\rm{5}}}} \right]{\rm{,}}\,\left[ {{\rm{Cr}}{{\left( {{\rm{CO}}} \right)}_{\rm{6}}}} \right]{\rm{,}}\,\left[ {{\rm{Ni}}{{\left( {{\rm{CO}}} \right)}_{\rm{4}}}} \right]\) etc. that contain an \({\rm{M}} – {\rm{CO}}\) bond are some of the examples of metal carbonyls.

Q.2: 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.

Q.3: Which carbonyl has the strongest \({\rm{CO}}\) Bond?
Ans: The metal carbonyls having the weakest \({\rm{M}} – {\rm{C}}\) bond has the strongest \({\rm{CO}}\) bond.

Q.4: What is the synergic effect in metal carbonyls?
Ans: The metal-carbon bond in metal carbonyls possesses both \({\rm{s}}\) and \({\rm{p}}\) character. The overlapping of the empty hybrid orbital of a metal atom with the filled hybrid orbital \(\left( {{\rm{HOMO}}} \right)\) of the carbon atom of carbon monoxide molecule results in the formation of an \(\left( {{\rm{M}} \leftarrow {\rm{CO}}} \right)\,{\rm{\sigma }} – \)bond. And the overlapping of filled dπ orbitals or hybrid dpπ orbitals of the metal atom with low-lying empty \(\left( {{\rm{LUMO}}} \right)\) orbitals of \({\rm{CO}}\) molecule results in the formation of \(\left( {{\rm{M}} \to {\rm{CO}}} \right)\,{\rm{\pi }}\) bonding. The metal to ligand bonding creates a synergic effect which strengthens the bond between \({\rm{CO}}\) and the metal.

Q.5: List down the uses of metal carbonyls.
Ans: The uses of metal carbonyls are as follows:
1. Metal carbonyls like carbonyl of iron are utilised for the preparation of inductors, pigments as dietary supplements.
2. Metal carbonyls are utilised as a catalyst in the production of a number of industrially essential compounds.