Chemistry > Coordination Compounds > 5.0 Valence bond theory
Coordination Compounds
1.0 Basics
2.0 Addition Salt
3.0 Nomenclature of Co-ordination Compounds
4.0 Werner's Co-ordination Theory
5.0 Valence bond theory
6.0 Crystal field splitting theory (CFST)
7.0 Effective atomic number
8.0 Magnetic Moment
9.0 Application of Crystal Field Splitting Theory (CFST)
10.0 Isomerism in Co-ordination compounds
11.0 Organo-metallic compounds
12.0 Stability of Co-ordination compounds
5.1 Postulates of Valence Bond Theory (VBT)
- Weak ligand as $O{H^ - }$ and ${H_2}O$ do not affect electronic configuration of central metal atom/ion.
- Strong ligand like $C{N^ - }$ and $CO$ form possible pair of unpaired electron present in central metal atom/ion.
- Complex compound having no unpaired electron are diamagnetic and complex compound having unpaired electron are paramagnetic.
- Magnetic behaviour of complex compound is given by $\mu = \sqrt {N(N + 2)} {\text{ BM}}$ where $N$ is number of unpaired electrons.
| S.No. | Co-ordination number | Hybridization | Orbital involved in hybridization | Structure |
| (i) | $2$ | $sp$ | $ns$, $np$ | Linear |
| (ii) | $3$ | $s{p^2}$ | $ns$, $n{p_x},n{p_y}$ | Trigonal planar |
| (iii) | $4$ | $s{p^3}$ $ds{p^2}$ | $ns,n{p_x},n{p_y},n{p_z}$ $(n - 1){d_{{x^2} - {y^2}}}$,$ns,n{p_x},n{p_y}$ | Tetrahedral Square planar |
| (iv) | $5$ | $s{p^3}$ $ds{p^3}$ | $(n - 1){d_{{z^2}}}$,$ns,n{p_x},n{p_y},n{p_z}$ $ns,n{p_x},n{p_y},n{p_z}$,$(n - 1){d_{{x^2} - {y^2}}}$ | Trigonal bipyramidal Square pyramidal |
| (v) | $6$ | ${d^2}s{p^3}$ $s{p^3}{d^2}$ | $ns,n{p_x},n{p_y},n{p_z},(n - 1){d_{{x^2} - {y^2}}},(n - 1){d_{{z^2}}}$ $ns,n{p_x},n{p_y},n{p_z},(n - 1){d_{{x^2} - {y^2}}},(n - 1){d_{{z^2}}},n{d_{{z^2}}}$ | Octahedral |
Inner orbital complex: If the complex is formed by use of inner d-orbital for hybridisation ${d^2}s{p^3}$. It is inner orbital complex. In the formation of inner orbital complex, the electron of the metal are forced pair up and hence the complex will be either dia-magnetic or will have lesser number of unpaired electron. Such a complex is called low spin complex.
Complexes with hybridisation $ds{p^2}$, $ds{p^3}$, ${d^2}s{p^3}$ are low spin complexes because they use inner vacant $d$ orbitals for hybridisation. This is only possible with the strong field ligands.
Example 1. ${\left[ {Ni(CN{}_4)} \right]^{ - 2}}$
Solution: Under the influence of the $C{N^ - }$ strong ligand as $N{H_3}$, the unpaired electrons get paired, inner complexes are formed by hybridisation $ds{p^2}$
1. Oxidation state =$x+4(-1)=-2$
2. $x$= $+2$
3. Hybridisation = Square planar
4. Nature = diamagnetic
Tip to remember: Complex compound in which metal cation has electronic configuration ${d^1}$ to ${d^3}$ always form inner orbital complex having hybridisation ${d^2}s{p^3}$ irrespective of nature of ligand.
Outer orbital complex: If a complex is formed by the use of outer d-orbitals for hybridisation ($s{p^3}{d^2}$), it is called outer orbital complex. The outer orbital complex will have large number of unpaired electron. Such a complex is also called High spin complex, it is always a Para magnetic nature.
Example 2. ${\left[ {Fe{{({H_2}O)}_6}} \right]^{ + 3}}$
Solution: Under the influence of ${{H_2}O}$, the unpaired electrons do not pair up and outer complex is formed by week ligands such as ${{H_2}O}$ ,${F^ - }$
1. Oxidation number = $x+6(0) = +3$
2. $x$ = $+3$
3. Hybridisation = ($s{p^3}{d^2}$)
4. Nature = paramagnetic
- It usually possible to predict the geometry of a complex from the knowledge of the magnetic behaviour on the basis of the valence bond theory.
| S.No. | Inner orbital Complex | Outer orbital Complex |
| 1. | Strong field or low spin ligands | Weak field or high spin ligands |
| 2. | Hybridisation is $ds{p^2}$ (where one orbital of $3d$, one orbital of $4s$ and two orbitals of $4p$) | Hybridisation is $s{p^3}$ (where one orbital of $4s$ and three orbital of $4p$) |
| 3. | Square planar shape | Tetrahedral complex |
Example 3. In the diamagnetic octahedral complex ${\left[ {Co{{(N{H_3})}_6}} \right]^{3 + }}$ the cobalt is in $+3$ oxidation state and has the electronic configuration $3d$. The hybridisation scheme is shown in the figure.
Here $N{H_3}$ is a strong ligand so electronic configuration done against hunds rule. Six pair of electrons, one from each $N{H_3}$ molecule, occupy the six hybrid orbital. Thus, the complex has octahedral geometry (fig) and is diamagnetic because of absence of unpaired electron. In the formation of this complex, in the inner $d$ orbital ($3d$) is used in hybridisation. The complex is called an Inner orbital Complex and Low spin Complex, where n=0, $\mu $=0 and hybridisation is ${d^2}s{p^3}$.
Example 4. In the paramagnetic octahedral complex ${\left[ {Co{F_6}} \right]^{3 - }}$ the cobalt is in +3 oxidation state and has the electronic configuration $3d$. The hybridisation scheme is shown in the figure.
Explanation: Here ${F^ - }$ is week ligand so electronic configuration done with according to hund's rule. Six pair of electrons, one from each ${F^ - }$ molecule, occupy the six hybrid orbital. Thus, the complex has octahedral geometry (fig) and is paramagnetic because of paired of electron. In the formation of this complex, since in the $d$ orbital ($4d$) is used in hybridisation ($s{p^3}{d^2}$). It is thus called Outer orbital Complex or High spin complex.
Example 5. In the square planar complexes, the hybridisation involed is $ds{p^2}$. An example is ${\left[ {Ni(CN{}_4)} \right]^{ - 2}}$. Here nickel is in $+2$ oxidation state and has the electronic configuration $3{d^6}$. The hybridisation scheme is shown in the figure.
Explanation: Each of the hybridised orbitals recieves a pair of electrons from a cyanide ion. The compound is DIAMAGNETIC as evident from the absence of unpaired electron.
- It is important to note that the hybrid do not actually exist. In fact, hybridisation is a Mathematical manipulation of wave form for the atomic orbitals involved.
$s{p^3}$ hybridisation
| $Ni{(CO)_4}$ | ${\left[{NiC{l_4}} \right]^{ - 2}}$ | ${\left[ {CuC{l_4}} \right]^{ - 2}}$ | ${\left[ {Zn{{(N{H_3})}_4}} \right]^{ + 2}}$ |
| $\left[ {MnC{l_4}} \right]{}^{ - 2},$ | ${\left[ {MnB{r_4}} \right]^{2 - }}$ | ${\left[ {Cu(CN){}_4} \right]^{ - 3}}$ | ${\left[ {CoC{l_4}} \right]^{2 - }}$ |
$ds{p^2}$ hybridisation
| ${\left[ {PtC{l_4}} \right]^{2 - }}$ | ${\left[ {Pt(C{N_4})} \right]^{4 - }}$ | ${\left[ {Ni{{(CN)}_4}} \right]^{2 - }}$ |
${d^2}s{p^3}$ hybridisation
| ${\left[ {Fe{{(CN)}_6}} \right]^{4 - }}$ | ${\left[ {Fe{{(CN)}_6}} \right]^{3 - }},$ | ${\left[ {Cr{{(N{H_3})}_6}} \right]^{3 + }}$ | ${\left[ {Co{{(OX)}_3}} \right]^{3 - }}$ | ${\left[ {Mn{{(CN)}_6}} \right]^{4 - }}$ | ${\left[ {Co{{(N{O_2})}_6}} \right]^{3 - }}$ |
| ${\left[ {Sc{F_6}} \right]^{3 - }}$ | ${\left[ {Ti{F_6}} \right]^{3 - }}$ | ${\left[ {Ti{{({H_2}O)}_6}} \right]^{3 + }}$ | ${\left[ {V{{({H_2}O)}_6}} \right]^{3 + }}$ | ${\left[ {Cr{{({H_2}O)}_6}} \right]^{3 + }}$ | $Cr{(CO)_6}$ |
$s{p^3}{d^2}$ hybridisation
| ${\left[ {Co{F_6}} \right]^{3 - }}$ | ${\left[ {Fe{{({H_2}O)}_6}} \right]^{3 + }}$ | ${\left[ {Ni(N{H_6})} \right]^{3 + }}$ | ${\left[ {Co{{({H_2}O)}_6}} \right]^{2 + }}$ | ${\left[ {Cr{{(N{H_3})}_6}} \right]^{2 + }}$ | ${\left[ {Fe{{(N{H_3})}_6}} \right]^{2 + }}$ |
Important points to remember:
1. In $C{o^{3 + }}$ complexes all the ligands acts as strong ligands except ${F^ - }$ .
2. $N{H_3}$ act as weak field ligand in case of $C{r^{2 + }},F{e^{2 + }}$ .
3. $P{t^{2 + }}, A{u^{2 + }}, P{d^{2 + }}$ preferrably form square planar complexes whether the ligand is strong or weak .
4. In ${d^1},{d^2}, {d^3}$ configuration, no pairing of electron takes place whether the ligand is strong or weak .
5. Weak ligand of $O{H^ - }$ , ${H_2}O$ do not affect the electronic configuration of central metal atom .
