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Valence Bond Theory in Coordination Compounds

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  • Last Updated : 14 Mar, 2022

There have been numerous approaches proposed to explain the nature of bonding in coordination molecules. One of them is the Valence Bond (VB) Theory. The Valence Bond Theory was developed to describe chemical bonding using the quantum mechanics method. This theory is largely concerned with the production of individual bonds from the atomic orbitals of the atoms involved in the formation of a molecule.

The electrons in a molecule, according to the valence bond hypothesis, occupy atomic orbitals rather than molecular orbitals. On bond formation, the atomic orbitals overlap, and the more the overlap, the stronger the bond. Metal bonding is mostly covalent in nature, and metallic structure involves the resonance of electron-pair connections between each atom and its neighbours.

The valence bond theory also explains the electronic structure of molecules created by the overlapping of atomic orbitals. It also stresses the fact that the nucleus of one atom in a molecule is drawn to the electrons of the other atoms.

Postulates of Valence Bond Theory

The following are the key postulates of valence bond theory.

  1. When two valence orbitals (half-filled) from separate atoms overlap on one other, covalent bonds develop. As a result of this overlapping, the electron density in the area between the two bonding atoms increases, boosting the stability of the resulting molecule.
  2. An atom’s valence shell has several unpaired electrons, allowing it to make many bonds with other atoms. According to the valence bond theory, the paired electrons in the valence shell do not participate in the creation of chemical bonds.
  3. Covalent chemical bonds are directed and parallel to the region corresponding to the overlapping atomic orbitals.
  4. Sigma bonds and pi bonds differ in the manner in which the atomic orbitals overlap, i.e. pi bonds are produced by sidewise overlapping, whereas sigma bonds are created by overlapping along the axis containing the nuclei of the two atoms.

Hybridization and Geometry of Complexes

Linus Pauling created the valence bond theory (VBT). His basic premise was that metal-ligand connections are created by the ligand contributing an electron pair to the metal, resulting in a coordinate bond between the metal and ligand.

Coordination Number

Types of Hybridization








Triangular planar[HgI3]



Tetrahedral [CoCl4]2-



Square planar[Ni(CN)4]2-



Tetrahedral MnO4



Trigonal bipyramidalFe(CO)5



Square pyramidal[Ni(CN)5]3-







Examples of Octahedral complexes

  • Inner Orbital Complexes: [Co(CN)6]3- ion
    • The oxidation state of cobalt in this combination is +3.
    • The electrical configuration of the Co+3 valence shell is 3d6.
    • Because the CN ligands are strong, they cause the pairing of 3d-electrons.
    • As a result, all six 3d electrons are coupled and occupy three of the five 3d orbitals.
    • Six d2sp3-hybrid orbitals are formed when the vacant 3d- orbitals combine with the vacant 4s and 4p orbitals.
    • Six coordinate covalent bonds are formed when these six hybrid orbitals overlap with six filled orbitals of ligands.
  • Outer Orbital Complexes: [FeF6]3- ion
    • The oxidation state of iron(Fe) in this combination is +3.
    • The electronic configuration of the Fe+3 valence shell is 3d5.
    • Because the F ligands are weak, no 3d-electrons are paired.
    • As a result, all five 3d electrons are occupied on five 3d orbitals.
    • The unoccupied 4s- orbitals interact with the vacant 4p orbitals and two vacant 5d orbitals to generate six sp3d2-hybrid orbitals.
    • Six coordinate covalent bonds are formed when these six hybrid orbitals overlap with six filled orbitals of ligands.

Number of Orbitals and Types of Hybridization

A metal atom or ion can utilize its (n-1)d, ns, np, or ns, np, nd orbitals for hybridization under the effect of ligands, generating a set of equivalent orbitals of definite geometry such as octahedral, tetrahedral, square planar, and so on. These hybrid orbitals can collide with ligand orbitals, allowing electron pairs to be donated for bonding.

Coordination Number

Type of Hybridization

Distribution of Hybrid Orbitals in Space




dsp2Square planar


sp3dTrigonal bipyramidal





Applications of Valence Bond Theory

  1. The valence bond theory’s greatest overlap criterion helps explain the creation of covalent bonds in numerous compounds.
  2. This is one of its most essential uses. The variation in the length and strength of chemical bonds in H2 and F2 molecules, for example, can be explained by differences in their overlapping orbitals.
  3. The covalent link in an HF molecule is produced by the overlap of the hydrogen atom’s 1s orbital and the fluorine atom’s 2p orbital, as explained by the valence bond theory.

Coordination Number Examples

The coordination number of a crystalline solid is the number of atoms, ions, or molecules that a central atom/ion has as its closest neighbours in the crystalline solid or coordination compound. According to our observations, magnesium has a strong coordination number of 6 and a strong affinity for water or other oxygen-containing ligands. For example, the coordination numbers of Pt and Fe in the complex ions [PtCl6]2– and [Fe(H2O)6]2+ are 6 and 6, respectively. Pt and Fe are linked to six mono-dentate ligands, Cl and H2O, in this case. Consider the compound [Cr(NH3)2Cl2Br2]. Because the total number of atoms/ions/molecules linked to Cr is discovered to be 6, the core atom Cr has coordination number 6 once more. Because en (ethylenediamine) is a bidentate ligand, the coordination number Co in the complex ion [Co(en)3]3+ is 6.

The metal-ligand connections may not all be at the same distance in some complexes. In some circumstances, an alternative meaning of coordination number is employed, which includes atoms that are further apart than their nearest neighbours. Some metals have a wavy structure. The structures of many chemical compounds are twisted. Unlike sodium chloride, which has cubic close-packed chloride ions, arsenic anions have hexagonal close-packed chloride ions.

Magnetic Properties of Complexes

  • Paramagnetic complexes have unpaired electrons in the core transition metal ion.
  • Diamagnetic complexes have no unpaired electrons in the core transition metal ion.
  • The spin only formula is used to compute a complex’s magnetic moment.

M = √[n(n+2)] BM

where, BM = Bohr Magneton

The magnetic moment of complex compounds is determined by:

  • Hybridization type.
  • The state of oxidation of a central transition metal ion.
  • The number of unpaired electrons.

Limitations of Valence Bond Theory

  1. Failure to account for carbon’s tetravalency.
  2. There is no information provided on the energy of the electrons.
  3. The idea assumes that electrons are concentrated in specific locations.
  4. It does not provide a quantitative assessment of the thermodynamic or kinetic stabilities of coordination molecules.
  5. There is no distinction between ligands that are weak and those that are powerful.
  6. There is no reason for the colour of coordination compounds.
  7. Although it provides a subjectively satisfying pictorial representation of the complex, it does not provide a quantitative understanding of complex stability.
  8. It predicts no distortion in symmetrical complexes, however, it predicts distortion in all copper (II) and titanium (III) complexes.
  9. It does not explain why the electrons must be placed against Hund’s rule at times while the electronic configuration is not disrupted at others.
  10. In the absence of an energy provider, the theory sometimes requires electrons to be transferred from a lower energy level (Example 3d) to a higher energy level (4p).
  11. Electron spin resonance demonstrates that the electron in Cu(II) complexes is not at the 4p level and that the complex is planar.
  12. It is unable to explain why some complexes are more labile than others. Labile complexes are those in which one ligand can easily be replaced by another. In contrast, inert complexes are ones in which ligand displacement is gradual.

Valence Bond and Molecular Orbital Theories Comparison

The language of bonding and antibonding orbitals, as well as electron delocalization, is introduced by molecular orbital theory into chemistry. The theory is provided here as an alternative to valence bond theory, with a very different formulation. Both theories, however, entail approximations to the actual electrical structures of molecules and can be improved. Valence bond theory is improved by integrating substantial ionic-covalent resonance; molecular orbital theory is improved by allowing for a range of molecular orbital occupation schemes. As these two improvement strategies are pursued, the wave functions created by the two approaches converge on one another, and the electron distributions predicted by the two approaches become identical.

When the molecular feature of interest is identified with the properties of individual bonds, valence bond theory is commonly used. As a result, it is frequently used in organic chemistry, where molecular reactions are frequently addressed in terms of the properties of their functional groups. Small confined sections of a molecule (such as a double bond) or specific groups of atoms are examples of the latter. Molecular orbital theory is commonly used to describe features that are commonly described in terms of delocalization. These properties include the spectroscopic properties of molecules, in which electromagnetic radiation is used to excite an electron from one molecular orbital to another, and all atoms contribute to the shift in electron density that occurs as a result of the excitation.

Sample Questions

Question 1: What are the shortcomings of VBT?


The valence bond hypothesis fails to explain carbon’s tetravalency and also fails to provide insight into the energies associated with electrons. The theory also assumes that electrons are concentrated in specific places.

Question 2: What are the merits of the valence bond theory?


The maximum overlap condition provided by the VBT can be used to explain how covalent bonds form in many compounds. The theory can also shed light on the ionic nature of chemical bonding.

Question 3: Why a Need for Valence Bond Theory


Lewis describes the arrangement of molecules using systems. It did not, however, describe the formation of chemical bonds. Similarly, VSEPR theory predicts the shape of simple molecules. However, its use was quite limited. It also lacked the ability to explain the geometry of a complicated atom. As a result, scientists were forced to publish the theory of valence bonds in order to address and overcome these restrictions.

Question 4: How are sigma and pi bonds formed?


When the atomic orbitals involved in the bond overlap head-to-head, a sigma bond is created. Pi bonds, on the other hand, entail parallel overlapping atomic orbitals.

Question 5: What is the orbital overlap concept?


According to this hypothesis, a covalent connection established between atoms causes the overlap of orbitals belonging to atoms with opposite spins of electrons. The type of overlapping between the atomic orbitals impacts the molecular orbital’s stability.

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