Concept of colour in metal complexes :

(1) Origin of colours:
(2) d-d transition:
(3) Selection rule of d-d transition:
(4) Charge transfer:

Colour originated by charge transfer when electronic transition occurs from one part of the Complex to other part i.e such type are also called internal Redox reaction. Intensity of clour in such type transition is very high as they donot require to follow any selection rule.

(A) Charge transfer from ligands to metal:
Examples: MnO₄ ^-1 , MnO₄ ^-2 , CrO₄ ^-2, Cr₂O₇ ^-2 , [Fe(H₂O)₅(NO)]⁺² , Na₄[Fe(CN)₅(NOS)].

(B) Charge transfer from metal to ligands:
Examples: [Fe(CO)₅] , [Fe (pi-C₂H₅)₂],  [Cr (pi-C₆H₆)₂],

(C) Charge transfer from metal to metal:

Examples: 

Prussian Blue: Fe4[Fe(CN)₆]₃    

Turnbull Blue: Fe4[Fe(CN)₆]₃

Related Question:

(1) Although both [Mn(H2O)6]2+ and [FeF6]3- have a d5 configuration and high-spin complexes. But the dilute solutions of Mn2+ and Fe +3 complexes are therefore colorless. Why?

(2) Which of the Complex of the following pairs has the highest value of CFSE?

(3) Colour of Complexes due to charge transfer:

(4) Why violet colour of [Ti(H2O)6]Cl3 disapear (colourless) on heating heating ? 

(5) Why [Ni(CN)4]-2 is colourless while [Ni(H2O)4]-2 although both have +2 oxidation state and 3d*8 configuration ?

(6) Why [FeF6]3– is colourless whereas [CoF6]3– is coloured ? 

(7) Why Fe(CO)5 is colourless while Fe(bipy)(CO)3 is intensely purple in colour ?

(8) Why all the tetrahedral Complexes are high spin Complexes ?

 

Postulates of Werner’s theory:

In coordination complex central atom metals exert two types of valencies; The important postulates of Werner’s theory are as follows:

(1) The primary or ionizable valencies which are satisfied by negative ions and equal the oxidation state of the central metal atom.

(2) The Secondary or non ionizable valencies which can be satisfied by neutral or negative ions/groups (or sometimes by cationic species -discovered later).

(3) The secondary valencies equal the coordination number of central metal atom/ion. This number is fixed for a metal.

(4) The ions/groups bound by the secondary linkages have characteristic spatial arrangements corresponding to different co-ordination numbers. In the modern terminology, such spatial arrangements are called coordination polyhedra. The various possibilities are

Coordination Number = 2 Linear

Coordination Number = 3 Triangular

Coordination Number = 4 Tetrahedral or sq. planar

Coordination Number = 5 square pyramidal or TBP

Coordination Number = 6 Octahedral

(5) The secondary valencies are generally represented by solid lines while the primary valencies are represented by dashed lines and the ions which satisfy both primary and secondary valencies will be drawn with both solid and dashed lines.

For example the complex [CoCl(H₂O)₅] Cl₂ is represented as

Introduction of Werner’s theory:

The systematic study of coordination compounds was started by a very famous Swiss scientist Alfred Werner whose pioneering work opened an entirely new field of investigation in inorganic chemistry. He prepared and characterized a large number of coordination compounds and studied their physical, chemical and isomeric behaviour by simple experimental techniques. On the basis of these studies. Werner, in 1898, propounded his theory of coordination compounds. Which is later termed as Werner’s Theory of Coordinate Compounds. Due to this theory he is awarded by Nobel prize and he is also called the ‘Father of Coordination Chemistry’.

CRYSTAL FIELD SPLITTING IN OCTAHEDRALCOMPLEXES:

For convenience, let us assume that the six ligands are positioned symmetrically along the Cartesian axes, with the metal atom at the origin. As the ligands approach, first there is an increase in the energy of d orbitals to that of the free ion just as would be the case in a spherical filed. Next, the orbitals lying along the axes (dz²and dx²-y² d) get repelled more strongly than d_xy, d_yz and d_xz orbitals, which have lobes directed between the axes. The d_xy , d_yz , d_xz orbitals are lowered in energy relative to the average energy in the spherical crystalfiled. Thus, the degenerate set of d orbitals get split into two sets: the lower energy orbitals set, t₂g and the higher energy, eg set. The energy separation is denoted by del.oct (the subscript o is for octahedral.

Crystal field stabilisation energy (CFSE):

The difference in energy of eg and t₂g Orbitals are called crystal field stabilisation energy (CFSE):

Where m and n = are number of electrons in t₂g and eg orbitals respectively and del.oct is crystalfield splitting energy in octahedral Complexes.

l = represents the number of extra electron pair formed because of the ligands in comparison to normal degenerate configuration.

P= (Pairing energy) the energy required for electron pairing in a single orbital. The actual configuration of complex adopted is decided by the relative values of delta and P

Case (1): If del.oct is less than P 

We have so called weak field or high spin situation, the fourth electron entered one of the eg orbitals giving configuration (t₂g³ and eg¹) If now 5th electron is added to a weak field the configuration become  (t₂g³ and eg²).

Case (2): If del.oct  is more than P , we have the strong field , low spin situation and pairing will occur in the t₂g level with eg level remaining unoccupied in entities of d¹ and d⁶ ions .

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CRYSTAL FIELD SPLITTING IN TETRAHEDRAL COMPLEXES:

Tetahedral complex (sp³):

In a tetrahedral field : Consider a cube such that a metal atom or ion is situated at its centre of symmetry through which the axis of geometry are passing and joining the face centres of this cube. Therefore, lobes of eg orbitals will be directed towards the face centres but those of t₂g orbitals will be pointing towards edge centres. Now consider 4 monodentate ligands approaching the metal, the 4 alternate corners of this cube so as to make a tetrahedron.

Thus it is clear that t₂g orbitals are nearer to the ligands than the eg orbitals. Hence t2g orbitals will experience more repulsion than eg orbitals. Therefore, crystal field splitting will be reversed of octahedral field which can be shown as below.

In tetrahedral complexes none of the ligand is directly facing any orbital so the splitting is found to be small in comparison to octahedral complexes. For the same metal, the same ligands and metal-ligand distances, it can be shown that del.tetra = (4/9) del.oct. This may attributes to the following two reasons.

(1) There are only four ligands instead of six, so the ligand field is only two thirds the size; as the ligand field spliting is also the two thirds the size and

 (2) The direction of the orbitals does not concide with the direction of the ligands. This reduces the crystal field spliting by roughly further two third.

Consequently, the orbital splitting energies are not sufficiently large for forcing pairing and, therefore, low spin configurations are rarely observed.

FACTORS FAVOURING TETRAHEDRAL COMPLEXES:

Tetrahedral complexes are favoured by steric requirements, either simple electrostatic repulsion of charge ligands or vander wall's repulsions of large one. A valence bond (VB) point of view ascribed tetrahedral structure to sp3 hybridisation.

Tetrahedral complexes are thus generally favoured by large ligands like Cl^-, B^-, I^- and PPh₃ and metal ions of six types;

(1) Those with a noble gas configuration such as Be²⁺ (ns⁰);

(2) Those with pseudo noble gas configuration (n-1) d¹⁰ns⁰np⁰, such as Zn^2+, Cu⁺ and Ga^3+, and

(3) Those transition metal ions which do not strongly favour other structure by virtue of the CFSE, such as Co^2+, d⁷.

(4) Those transition metal which have lower oxidation state.

(5) Those metals generally with electronic configuration d⁰, d⁵ and d¹⁰ prefer to form such complexes.

(6) It is observed that

OTHER EXAMPLES :

SN	Complex	Nature
1	[Ni(CO)4]	Diamagnetic
2	[Ni(Cl)4]2-	Paramagnetic with two unpaired electron
3	[NiCl2(pph3)2]	Paramagnetic with two unpaired electron
4	[MnCl4]2-	Paramagnetic with five unpaired electron
5	[FeCl4]2-	Paramagnetic with four unpaired electron
6	[Cu(py)4]+	Diamagnetic
7	Cs2[CuCl4]	Paramagnetic with two unpaired electron (Orange tetrahedral) Sp3
8	NH3[CuCl4]	Paramagnetic with two unpaired electron (Yellow Square Planer) dsp2
9	[Zn(NH3)4]2+	(d10) CFSE=0 , Diamagnetic
10	[Zn(CN)4]2-	(d10) CFSE=0 , Diamagnetic