JEE Main & Advanced Chemistry Coordination Chemistry Bonding in co-ordination compounds (Werner's Coordination theory)

Werner was able to explain the bonding in complex.

Primary valency (Pv) : This is non- directional and ionizable. In fact it is the positive charge on the metal ion.

Secondary valency  (Sv) : This is directional and non- ionizable. It is equal to the number of ligand atoms co-ordinated to the metal (co-ordination number). Example :

Nature of the complex can be understood by treating the above complexes with excess of \[AgN{{O}_{3}}.\]

\[CoC{{l}_{3}}.\,6N{{H}_{3}}\to 3AgCl,\,\,[Co{{(N{{H}_{3}})}_{6}}C{{l}_{3}}\](three chloride ion)

\[CoC{{l}_{3}}.\,5N{{H}_{3}}\to 2AgCl,\,\,[Co{{(N{{H}_{3}})}_{5}}C{{l}_{2}}\](two chloride ion)

\[CoC{{l}_{3}}.\,4N{{H}_{3}}\to 1AgCl,\,\,[Co{{(N{{H}_{3}})}_{4}}C{{l}_{2}}\] (one chloride ion)

\[CoC{{l}_{3}}.\,3N{{H}_{3}}\to no\,AgCl,\,\,[Co{{(N{{H}_{3}})}_{3}}C{{l}_{3}}\](no  chloride ion)

The nature of bonding between central metal atom and ligands in the coordination sphere has been explained by the three well-known theories. These are :

(1) Valence Bond theory of coordination compounds

(i) The suitable number of atomic orbitals of central metal ion (s, p, d) hybridise to provide empty hybrid orbitals.

(ii) These hybrid orbitals accept lone pair of electrons from the ligands and are directed towards the ligand positions according to the geometry of the complex.

(iii) When inner d-orbitals i.e. \[(n-1)d\] orbitals are used in hybridization, the complex is called – inner orbital or spin or hyperligated complex.

(iv) A substance which do not contain any unpaired electron is not attracted by 2 magnet. It is said to be diamagnetic. On the other hand, a substance which contains one or more unpaired electrons in the electrons in the d-orbitals, is attracted by a magnetic field [exception \[{{O}_{2}}\] and NO]. It is said to be paramagnetic.

Paramagnetism can be calculated by the expression, \[{{\mu }_{s}}=\sqrt{n(n+2),}\] where \[\mu =\] magnetic moment.

s= spin only value and n= number of unpaired electrons.

Hence, if \[n=1,\,{{\mu }_{s}}=\sqrt{1(1+2)}=1.73\,B.M.\], if \[n=3,\,{{\mu }_{s}}\] \[=\sqrt{3(3+2)}=3.87\,B.M.\]and so on

On the basis of value of magnetic moment, we can predict the number of unpaired electrons present in the complex. If we know the number of unpaired electrons in the metal complex, then it is possible to predict the geometry of the complex species.

(v) There are two types of ligands namely strong field and weak field ligands. A strong field ligand is capable of forcing the electrons of the metal atom/ion to pair up (if required). Pairing is done only to the extent which is required to cause the hybridization possible for that co-ordination number. A weak field ligand is incapable of making the electrons of the metal atom/ ion to pair up.

Strong field ligands : \[C{{N}^{-}},CO,\,en,\,N{{H}_{3}},{{H}_{2}}O,\,N{{O}^{-}},Py\].

Weak field ligands :

\[{{I}^{-}},B{{r}^{-}},C{{l}^{-}},{{F}^{-}},NO_{3}^{-},O{{H}^{-}},{{C}_{2}}O_{4}^{2-},N{{H}_{3}},{{H}_{2}}O\].

Limitations of valence bond theory

The valence bond theory was fairly successful in explaining qualitatively the geometry and magnetic behaviour of the complexes. But, it could not explain the following :

 (i) The origin of their absorption spectra could not be explained.

 (ii) Why did different complexes of the same metal show different colours.

 (iii) Relative stabilities of different complexes could not be explained.

 (iv) Why should certain ligands form high spin, while others low spin complexes.