VALENCE BOND THEORY

      It is the simplest of the three theories and was developed mainly by Pauling. It describes the bonding in terms of hybridized orbitals of the central metal atom or ion. The theory mainly deals withe the geometry and magnetic properties of the complexes. The salient features of the theory are:

(i) The central metal loses a requisite number of electrons to form the ion. The number of electrons lost is the valency of the resulting cation. In some cases, the metal atom does not lose electrons.

(ii) The central metal ion or atom (as the case may be) makes available a number of empty s-, p- and d-atomic orbitals equal to its coordination number. These vacant orbitals hybridize together to form hybrid orbitals which are same in the number as the atomic orbitals hybridizing together. They are vacant, equivalent in energy and have definite geometry.

Some of the common hybridized orbitals met in the coordination compounds are listed below:

Coordination number	Type of hybridization	Geometry	Examples
2	sp	Linear	[Ag(NH3)2]+ ; [Ag(CN)2]-
3	sp2	Trigonal planar	[HgI3]-
4	sp3 dsp2	Tetrahedral Square planar	Ni(CO)4, [Ni(X4]2- [ZnCl4]2-, [CuX4]2- Where X= Cl-, B- , I- [Ni(CN)4]2- , [Cu(NH3)4]2- [Ni(NH3)4]2=
5	dsp3 sp3d	Trigonal bipyramidal Square pyramidal	Fe(CO)5, [CuCl5]3- [SbF5]2-
6	d2sp3 or sp3d2	Octahedral (Inner orbital) (Outer orbital)	[Cr(NH3)6]3+ ; [Fe(CN)6]3- [Fe(H2O)6]2+ ; [Ni(NH3)6]2+ [FeF6]3-

(iii) The non-bonding electrons of the metal occupy the inner orbitals. These ate grouped in accordance with Hund's rule, however, under the influence of some strong ligands, there may be some re-arrangement of electrons in the atomic orbitals the d-orbitals participating in this process of hybridization may be either (n-1)d2sp3 or nsp3d2. The complexes thus formed are referred to as inner or low spin and outer or high spin complexes, respectively.

(iv) The ligands have at least one o-orbital containing a lone pair of electrons. Vacant hybrid orbitals of the metal atom or ion overlap with the o-orbitals containing lone pair or electrons of the ligands to form M<- ligand o-bond.This bond is called coordinate bond and possesses a considerable amount of polarity.

(v) It is  possible to predict the magnetic properties of the complex if the geometry of the complex ion is known. If the complex contains unpaired electrons, it is paramagnetic in nature whereas if it does not contain unpaired electrons all re paired, the complex is diamagnetic in nature.
the number of unpaired electrons and the geometries of the complex ions having central metal ion with configurations d1 to d9 are related to each other as shown below:

dx Configuration	Number of unpaired electrons for different geometries
				Octahedral
				Inner orbitals (d2sp3)	Outer orbitals (sp3d2)
d1	1		1	1	1
d2	2		2	2	2
d3	3		3	3	3
d4	4		4	2	4
d5	5		3	1	5
d6	4		2	0	4
d7	3		1	1 (Shifted to higher orbits)	3
d8	2		0	0 (2 electrons shifted)	2
d9	1		1 (Shifted)	1 (3 electrons shifted)	1

           

        Limitation of valence bond theory:

The valence bond theory was fairly successful in explaining qualitatively the geometry and magnetic properties of complexes. However, it has a number of limitations.

(i) The theory does not offer any explanation about the spectra of complex (why most of the complexes are coloured).

(ii) Sometimes the same metal ion assumes different geometry when formation of complex ion takes place.The theory is unable to explain why at one time the electrons are rearranged against the Hund's rule while at other times the electronic configuration is not disturbed.

(iii) The theory does not offer an explanation for the existence of inner-orbital and outer-orbital complexes.

(iv) The theory does not explain why certain complexes are labile while others are inert.

(v) In the formation of [Cu(NH3)4]2+, one electron is shifted from 3d to 4p-orbital. The theory is silent about the energy availability for shifting such an electron. Such an electron can be easily lost, then why [Cu(NH3)4]2+ complex does not show reducing properties.

(vi) The changes in energies of the metal orbitals on formation of complex are difficult to calculate mathematically. The properties of the complexes are more satisfactorily explained by other two theories. The approach of ligands towards a metal or ion creates a field which is responsible for the splitting of d-orbitals into different energy levels. The extent of splitting depends on the nature and number of ligands which surround the metal atom or ion and explains the magnetic and spectroscopic properties of the complex. More details about these theories are dealt in higher classes.