Subject Chemistry Paper No and Title 11 : Inorganic Chemistry...

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PAPER No. 11 : Inorganic Chemistry-III Module No 4 : Molecular orbitals of CO and Nature of bonding in metal carbonyls

Subject Chemistry

Paper No and Title 11 : Inorganic Chemistry-III

Module No and Title 4 : Molecular orbitals of CO and Nature of bonding in metal carbonyls

Module Tag CHE_P11_M4

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TABLE OF CONTENTS 1. Learning Outcomes

2. Introduction

3. Molecular orbital theory

4. Molecular orbital theory applied to CO

5. Nature of bonding in linear M – C – O linkage

6. Nature of bonding in metal-metal (M – M) bond

7. Nature of bonding in bridging M – CO – M linkage

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1. Learning Outcomes

After studying this module, you shall be able to

• Know about nature of linear M-C-O linkage • Learn how the molecular orbital theory applied to CO • Analyze the different modes of attachment of CO to the metal ion.

2. Introduction

Lewis formula of CO: Carbon atom has four valence electrons while oxygen atom has six valence electrons. Atom Atomic no. Electronic configuration K L C 6 2 4 O 8 2 6 Carbon and oxygen atoms share three electron pairs between them to complete octet leading to the formation of triple covalent bond as follows:

Thus both carbon and oxygen atoms complete the octet of electrons in their valence shells. In two bond pairs carbon and oxygen atoms provide one electron each and thus two normal covalent bonds result. In the third bond pair, both the electrons are provided by oxygen atom and thus it leads to the formation of a coordinate bond. Both carbon and oxygen atoms of carbon monoxide molecule have one lone pair of electrons. Hybridisation approach of CO Both carbon and oxygen atoms of CO molecule are sp hybridised.

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The singly occupied spz hybrid orbital of carbon atom overlaps with the singly occupied spz hybrid orbital of oxygen atom in linear fashion forming σ-bond.

The singly occupied 2px atomic orbital of carbon atom overlaps with the singly occupied 2px atomic orbital of oxygen atom in sidewise manner forming a π-bond.

The filled 2py atomic orbital of oxygen atom overlaps with the vacant 2py atomic orbital of carbon atom in sidewise manner forming O C coordinate bond which is also a π-bond. According to this concept there are three covalent bonds between carbon and oxygen atoms in CO molecule and both carbon and oxygen atoms have lone pair of electrons in spz hybrid orbitals.

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3. Molecular orbital theory

The molecular orbital theory was put forward by Hund and Mulliken in order to explain the nature of covalent bond. The main points of this theory are given below:

i) The molecular orbitals are produced by the linear combination of atomic orbitals (LCAOs) of comparable energies lying on two or more atoms-same or different.

ii) The atomic orbitals taking part in linear combination should have the similar symmetry property about the internuclear axis.

iii) The number of molecular orbitals (MOs) produced is always equal to the number of atomic orbitals (AOs) taking part in linear combination.

iv) If two atomic orbitals take part in linear combination two molecular orbitals are produced. One of these molecular orbitals has lower energy than both the atomic orbitals. It is called bonding molecular orbital. The other molecular orbital which has higher energy than both the atomic orbitals is called antibonding molecular orbital.

Where, ψb = Bonding molecular orbital

Ψa = Antibonding molecular orbital

v) Bonding molecular orbital is formed by the addition of atomic orbitals while the antibonding molecular orbital is formed by the subtraction overlap of atomic orbitals.

Ψb = ΨX + ΨY

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Ψa = ΨX – ΨY

vi) If both the atomic orbitals have the same energy they contribute equally to bonding and antibonding molecular orbitals. If one of the atomic orbitals has lower energy than the other, it will contribute more to the bonding molecular orbital.

Ψb = xΨX + yΨY

y>x, if Y is more electronegative than X.

Conversely, more stable atomic orbital contributes less to the antibinding molecular orbital.

Ψa = xΨA - yΨB

vii) Aufbau principle, Hund’s rule and Pauli’s exclusion principle are applicable to the molecular orbitals also.

viii) If an electron enters bonding molecular orbital, the electron density between two nuclei increases. It increases the attractive force between two atoms and decreases the energy of the system. On the other hand, when an electron enters antibonding molecular orbital, the electron density between two nuclei decreases. It increases the repulsive force between two atoms and increases the energy of the system.

ix) Each electron pair of the bonding molecular orbital constitutes one covalent bond while each electron pair of the antibonding molecular orbital cancels one covalent bond.

x) The bond order, i.e., the number of bonds in a molecule is given by

Bond order (B.O.) =

Where, Nb = Number of electrons in the bonding molecular orbitals

Na = Number of electrons in the antibonding molecular orbitals

xi) Any molecule is stable only when the bond order is either equal to or greater than 1.

B.O. = 0; the molecule does not exist

0 < B.O. < 1; the molecule is unstable

4. Molecular orbital theory applied to CO:

The molecular orbital energy level diagram of CO molecule is given on the next page. Oxygen atom is more electronegative and smaller in size than carbon atom. Therefore, the orbitals of

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oxygen atom have lower energy than corresponding orbitals of carbon atom. The spz hybridization takes place on both carbon and oxygen atoms. Two hybrid orbitals of carbon or oxygen atom are not identical. The lower energy spz hybrid orbital has greater s-character(s-rich) and is more electronegative while the higher energy spz hybrid orbital has greater p-character (p-rich) and is less electronegative. The lower energy spz hybrid orbital of carbon atom is comparable in energy with the higher energy spz hybrid orbital of oxygen atom and they combine together forming bonding (σ) and antibonding (σ*) molecular orbitals. The very low-energy spz hybrid orbital of oxygen atom and the high-energy s-pz of carbon atom are far apart in energy and they remain non-bonding and they are localized on oxygen and carbon atoms respectively. These hybrid orbitals are occupied by lone pair of electrons and are directed away from CO bond. The bond order, i.e., the number of bonds in CO molecule is given by

Bond order (B.O.) =

=

= 3

(Molecular orbital energy level diagram of CO molecule)

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(i) Atomic orbitals of carbon atom before hybridization

(ii) Atomic orbitals /hybrid orbitals of carbon atom after hybridization

(iii) Atomic orbitals /hybrid orbitals of oxygen atom after hybridization

(iv) Atomic orbitals of oxygen atom before hybridization

The 2px atomic orbital of carbon atom and the 2px atomic orbital of oxygen atom take part in linear combination forming bonding(π2px) and antibonding ((π2px

*) molecular orbitals .Similarly, The 2py atomic orbital of carbon atom and the 2py atomic orbital of oxygen atom take part in linear combination forming bonding(π2py) and antibonding ((π2py

*) molecular orbitals. The σ bonding and π bonding MOs are polarized towards more electronegative oxygen atom while the antibonding π* MOs are polarized towards less electronegative carbon atom as shown in the following figure.

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The molecular parameters of CO molecule are given below:

Bond order = 3

Bond length = 1.131 Å

Bond energy = 1072 kilojoules per mole

Magnetic property = Diamagnetic

IR absorption peak at 2143 cm-1

According to molecular orbital theory, the carbon monoxide molecule can be represented as:

p-rich orbital s-rich orbital

less electronegative more electronegative

more directional less directional

5. Nature of bonding in linear M – C – O linkage:

CO molecule is an ambidentate ligand. It has lone pair of electrons on both carbon and oxygen atoms. But the orbital containing the lone pair of electrons on carbon atom is p-rich, larger in size, less electronegative and more directional while the orbital containing lone pair of electrons in oxygen atom is s-rich, smaller in size, more electronegative and less directional. Therefore CO ligand is coordinated to metal atom in metal carbonyls invariably through its carbon atom and through oxygen atom. The mononuclear metal carbonyls contain only linear M – C – O linkage. Let us illustrate the nature of bonding in the linear M – C – O linkage taking the example of Cr(CO)6.This complex has octahedral shape arising from d2sp3 hybridisaton on chromium atom.

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The filled spz hybrid orbital of carbon atom of CO molecule overlaps with the vacant d2sp3 hybrid orbital of Cr atom forming Cr CO σ- bond. Thus the Cr –C – O linkage is linear. Six such bonds are formed in Cr(CO)6 molecule.

Vacant d2sp3 hybrid orbital Filled spz hybrid orbital Cr CO σ- bond

The filled 3dxy , 3dyz or 3dxz atomic orbital of Cr atom(dπ orbitals) overlaps with the vacant π*(π2px or π2py ) molecular orbital of CO molecule forming Cr CO π- bond which is also known as back –bond.Cr atom is capable of forming three such bonds at a time with three CO ligands using its three dπ orbitals.

Filled 3dxy, 3dyz or3dxz orbital Vacant π* MO Cr CO π bond

An average bond order equal to 1.5(1σ + is established in Cr – CO bond due to resonance.

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CO molecule is poor σ-donor. But the formation of Cr CO π- bond increases the electron density on CO molecule and thus increases the σ-donating ability of CO molecule. Thus the formation of M CO π- bond is very important in determining the stability of metal carbonyls. It not only increases the M –CO bond order but it also strengthens the M CO σ-bond. This type of bonding in which two bonds reinforce each other is called synergic bonding.

Any metal carbonyl is stable only when M CO π-bond of significant strength is formed. For example, [Ni(CO)4] is a stable metal carbonyl while the isoelectronic [Zn(CO)4]2- does not exist due to inability of Zn2+ to form Zn2+ CO, CO π-bond on account of its dispositive charge.

6. Nature of bonding in metal –metal (M –M) bond:

The metal-metal (M – M) bond in binuclear and polynuclear metal carbonyls may be formed by overlapping of singly occupied hybrid orbitals or d atomic orbitals of two metal atoms. The M – M bond may be a normal σ-bond or bent σ- bond or a δ- bond. In Mn2(CO)10 , the Mn – Mn bond is formed by the linear overlapping of singly occupied d2sp3 hybrid orbitals of two Mn atoms.

In Co2(CO)8 , the Co – Co bond is formed by the bent overlapping of singly occupied d2sp3 hybrid orbitals of two Co atoms.

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Resulting in the formation of bent Co – Co bond In Fe2(CO)9 , the Fe – Fe bond is formed by the lateral overlapping of singly occupied 3dxy

atomic orbitals of two Fe atoms. It is a δ-bond involving four overlap zones.

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7. Nature of bonding in bridging M – CO – M linkage:

CO ligand has a tendency to bridge two metals. In the most cases, the M – M bond is accompanied with the CO bridging. The carbon atom of the bridging carbonyl group may be regarded as sp2 hybridised. The carbon atom of the bridging carbonyl group forms two σ-‐bonds with neighboring two metal atoms. This bond is formed by the overlapping of the singly occupied sp2hybrid orbital of the carbon atom and singly occupied hybrid orbital of the metal atom. Thus the bridging carbonyl group forms the bent linkage.

M M

C

O

The carbon-‐oxygen bond order in the bridging carbonyl group is almost equal to 2.In Co2(CO)8 , the bonding in the bridging Co – CO – Co linkage is shown in the following figure.

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Bent Co – Co bond singly occupied sp2 hybrid orbital singly occupied d2sp3 hybrid orbital

The ability of the carbonyl group to act as the bridging ligand decreases with increase in the M – M internuclear distance as expected as evident from the following example.

Metal carbonyl Central metal Atomic radius of metal No. of bridging carbonyl groups

Fe2(CO)9 Fe 1.17Å 3

Os2(CO)9 Os 1.26Å 1

Fe3(CO)12 Fe 1.17Å 2

Os3(CO)12 Os 1.26Å 0

8. Summary

• Both carbon and oxygen atoms of CO molecule are sp hybridized with C- O bond order 3 • CO molecule has one electron pair on both carbon and oxygen atoms • The hybrid orbital containing the lone pair on carbon atom is p-rich, larger in size and

more directional • The hybrid orbital containing the lone pair on oxygen atom is s-rich, smaller in size and

less directional • CO ligand is attached to the metal atom in metal carbonyls invariably through its carbon

atom • The terminal carbonyl group has the linear M – C – O linkage

CO

• The bridging carbonyl group has the bent M M linkage

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• CO is a σ-donor as well as a π-acceptor ligand • The filled sp hybrid orbital of carbon of CO molecule overlaps with the vacant hybrid

orbital of the metal atom forming the M CO σ-bond.

Subject Chemistry Paper No and Title 11 : Inorganic Chemistry...

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