Notes-Part-2-Class-12-Chemistry-Chapter-9-Coordination Compounds-Maharashtra Board

Topics to be Learn : Part-1 Introduction Types of ligands Terms used in coordination chemistry Classification of complexes IUPAC nomenclature of coordination compounds Effective Atomic Number (EAN) Rule Isomerism in coordination compounds Topics to be Learn : Part-2 Stability of the coordination compounds Theories of bonding in complexes Applications of coordination compounds

Stability of the coordination compounds:

• The stability of coordination compounds can be explained by knowing their stability constants.
• The stability is governed by metal-ligand interactions.
• In this the metal serves as electron-pair acceptor while the ligand as Lewis base (since it is electron donor).
• The metal-ligand interaction can be realized as the Lewis acid-Lewis base interaction. Stronger the interaction greater is stability of the complex.

Consider the equilibrium for the metal ligand interaction :

M^a+ + nL^x— ⇌ [MLn]^(a+)+(nx—)

where a, x, [a+ + nx- ] denote the charge on the metal, ligand and the complex, respectively.

Now, the equilibrium constant K is given by

K = \(\frac{[ML_n]^{(a+)+(nx-)}}{[M^{a+}][L^{x-}]_n}\)

Stability of the complex can be explained in terms of K. Higher the value of K larger is the thermodynamic stability of the complex.

The equilibria for the complex formation with the corresponding K values are given below.

Ag⁺ + 2CN ⇌  [Ag(CN)₂] ^— K = 5.5 ×10¹⁸

Cu²⁺ + 4CN^— ⇌  [Cu(CN)₄]^2— K = 2.0 ×10²⁷

Co³⁺ + 6NH₃ ⇌ [Co(NH₃)₆]³⁺ K = 5.0 ×10³³

From the above data, [Co(NH₃)₆]³⁺ is more stable than [Ag(CN)₂] ^— and [Cu(CN)₄]^2—.

Theories of bonding in complexes :

The metal-ligand bonding in coordination compounds has been described by Valence bond theory (VBT) and Crystal field theory (CFT).

Valence bond theory (VBT) :

• Vacant d-orbitals of metal ion form coordination bonds with ligands.
• s, p orbitals along with vacant d—orbitals of metal ion take part in hybridisation.
• The number of vacant hybrid orbitals formed is equal to number of hybridising orbitals which is equal to the number of ligand donor atoms or coordination number of the metal.
• The metal-ligand coordination bonds are formed by the overlap between the vacant hybrid orbitals of metal and the filled orbitals of the ligands.
• The hybrid orbitals used by the metal ion point in the direction of the ligands.
• When inner (n — 1)d orbitals of metal ion are used in the hybridisation then the complex is called (a) inner orbital complex while when outer nd orbitals are used, complexes are called (b) outer orbital complexes.

Steps to understand the metal-ligand bonding :

• Find oxidation state of central metal ion
• Write valence shell electronic configuration of metal ion.
• See whether the complex is low spin or high spin. (applicable only for octahedral complexes with d⁴ to d⁸ electronic configurations).
• From the number of ligands find the number of metal ion orbitals required for bonding.
• Identify the orbitals of metal ion available for hybridisation and the type of hybridisation involved.
• Write the electronic configuration after hybridisation.
• vii Show filling of orbitals after complex formation.
• viii.Determine the number of unpaired electrons and predict magnetic behavior of the complex.

Salient features of valence bond theory (VBT) :

• According to this theory, a central metal atom or ion present in a complex provides a definite number of vacant orbitals (s, p, d and f ) to accommodate the electrons from the ligands for the formation coordinate bonds with the metal ion / atom.
• The number of vacant orbitals provided by the central metal atom or ion is the same as the coordination number of the metal. For example : Cu²⁺ provides 4 vacant orbitals in the complex. [Cu(NH₃)₄] ²⁺.
• The vacant orbitals of metal atom or ion undergo hybridisation forming the same number of hybridised orbitals, since the bonding with the hybrid orbitals is stronger.
• Each ligand has one or more orbitals containing one or more lone pairs of the electrons.
• The shape or geometry of the complex depends upon the type of hybridisation of the metal ion / atom.
• When inner orbitals namely (n — 1) d orbitals in transition metal atom or ion hybridise, the complex is called inner complex and when outer orbitals i.e., nd orbitals hybridise then the complex is called outer complex.
• When the central metal atom or ion in the complex contains one or more unpaired electrons the complex is paramagnetic while if all the electrons are paired, the complex is diamagnetic.

Octahedral, complexes :

Structure of octahedral complex, [CO(NH₃)₆]³⁺ on the basis of valence bond theory : (low spin)

(i) Hexaamminecobalt(III) ion, [CO(NH₃)₆]³⁺ is a cationic complex, the oxidation state of cobalt is +3 and the coordination number is 6.

(ii) Electronic configuration : ₂₇Co [Ar]¹⁸ 3d⁷ 4s²

      Electronic configuration : Co³⁺ [Ar]¹⁸ 3d⁶ 4s⁰ 4p⁰

(iii) Since NH₃ is a strong ligand, due to spin pairing effect, all the four unpaired electrons in 3d orbital are paired giving two vacant 3d orbitals.

(iv) Since the coordination number is 6, Co³⁺ ion gets six vacant orbitals by hybridisation of two 3d vacant orbitals, one 4s and three 4p orbitals forming six d²sp³ hybrid orbitals giving octahedral geometry It is an inner complex.

(v) 6 lone pairs of electrons from 6NH₃ ligands are accommodated in the six vacant d²sp³ hybrid orbitals. Thus six hybrid orbitals of Co³⁺ overlap with filled orbitals of NH₃ forming 6 coordinate bonds giving octahedral geometry to the complex.

(vi) As all electrons are paired the complex is diamagnetic.

Geometry of [CoF₆]^3— (high spin) on the basis of valence bond theory :

(i) Oxidation state of central metal Co is 3+

(ii) Valence shell electronic configuration of Co³⁺ is

(iii) Six fluoride F ligands, thus the number of vacant metal ion orbitals required for bonding with ligands would be six.

(iv) Complex is high spin, that means pairing of electrons will not take place prior to

hybridisation. Electronic configuration would remain the same as in the free state shown above.

(v) Six orbitals available for the hybridisation. Those are one 4s, three 4p, two of 4d orbitals

Six metal orbitals after bonding with six F ligands led to the sp³d² hybridization. The d orbitals participating in hybridisation for this complex are nd.

(vi) Six vacant sp³d² hybrid orbital of Co³⁺ overlap with six orbitals of fluoride forming Co - F coordinate bonds.

(vii) Configuration after complex formation.

(viii) The complex is octahedral and has four unpaired electrons and hence, is paramagnetic.

Tetrahedral complex :

Structure of tetrachloronickelate(II) [Ni(Cl)₄]^2—on the basis of valence bond theory :

(i) Oxidation state of nickel is +2

(ii) Valence shell electronic configuration of Ni²⁺

(iii) number of Cl^— ligands is 4. Therefore number of vacant metal ion orbitals required for bonding with ligands must be four.

(iv) Four orbitals on metal available for hybridisation are one 4s, three 4p. The complex is tetrahedral.

The four metal ion orbitals for bonding with Cl^— ligands are derived from the sp³ hybridization.

(vi) Four vacant sp³ hybrid orbital of Ni²⁺ overlap with four orbitals of Cl^— ions.

(vii) Configuration after complex formation would be.

(viii)The complex has four unpaired electrons and hence, paramagnetic.

Magnetic moment μ is, μ = \(\sqrt{n(n+2)}=\sqrt{2(2+2)}\)  = 2.83 B.M

Square planar complex :

Structure of  [Ni(CN)₄]^2—on the basis of valence bond theory :

(i) Oxidation state of nickel is +2

(ii) Valence shell electronic configuration of Ni²⁺

(iii) Number of CN ligands is 4, so number of vacant metal ion orbitals required for bonding with ligands would be four.

(iv) Complex is square planar so Ni²⁺ ion uses dsp² hybrid orbitals.

(v) 3d electrons are paired prior to the hybridisation and electronic configuration of Ni²⁺ becomes :

(vi) Orbitals available for hybridisation are one 3d, one 4s and two 4p which give dsp² hybridization.

(vii) Four vacant dsp2 hybrid orbitals of Ni²⁺ overlap with four orbitals of CN ions to form Ni - CN coordinate bonds.

(vii) Configuration after the complex formation becomes.

(viii) The complex has no unpaired electrons and hence, dimagnetic.

Remember :

Coordination number	Geometry of complex	Hybridization
2	Linear	sp
4	Tetrahedral	sp3
4	Square planar	dsp2
6	Octahedral	d2sp3/ sp3d2

Crystal Field Theory (CFT) :

Assumptions of Crystal Field Theory (CFT) :

Bethe and van Vleck developed Crystal Field Theory (CFT) to explain various properties of coordination compounds. The salient features of CFT are as follows :

In a complex, the central metal atom or ion is surrounded by various ligands which are either negatively charged ions (F^—, Cl^—, CN^—, etc.) or neutral molecules (H₂O, NH₃, en, etc.) and the most electronegative atom in them points towards central metal ion. .

The ligands are treated as point charges involving purely electrostatic attraction between them and metal ion.

• The central metal ion has five, (n — 1)d degenerate orbitals namely d_xy, d_yz, d_zx, d_(x² _{– y}²₎ and d_z².
• When the ligands approach the metal ion, due to repulsive forces, the degeneracy of d-orbitals is destroyed and they split into two groups of different energy, t_2g and e_g orbitals. This effect is called crystal field splitting which depends upon the geometry of the complex.
• The d-orbitals lying in the direction of ligands are affected to a greater extent while those lying in between the ligands are affected to a less extent.
• Due to repulsion, the orbitals along the axes of ligands acquire higher energy while those lying in between the ligands acquire less energy.
• Hence repulsion by ligands give two sets of split up orbitals of metal ion with different energies.

Crystal field Splitting in an octahedral complex

• The energy difference between two sets of d-orbitals after splitting by ligands is called crystal field splitting energy (CFSE) and represented by Δ₀ or by arbitrary term 10D_q. The value of Δ or 10D_q depends upon the geometry of the complex.

The electrons of metal ion occupy the split d-orbitals according to Hund’s rule, aufbau principle and those orbitalswith minimum repulsion and the farthest away from the ligands.

CFT does not account for overlapping of orbitals of central metal ion and ligands, hence does not consider covalent nature of the complex.

From the crystal field stability energy, the stability of the complexes can be known.

Crystal field splitting :

The splitting of five degenerate d-orbitals of the transition metal ion into different sets of orbitals (t_2g and e_g)having different energies in the presence of ligands in the complex is called crystal field splitting.

Crystal field stabilisation energy ((CFSE) : It is the change in energy achieved by preferential filling up of the orbitals by electrons in the complex of metal atom or ion.

• CFSE is expressed as a negative quantity i.e., CFSE ≤ 0. Higher the negative value more is the stability of the complex.

Factors affecting Crystal Field Splitting parameter (Δ₀)

Crystal Field Splitting parameter (Δ₀) depends on, (a) Strength of the ligands and (b) Oxidation state of the metal.

(i) Strength of the ligands : The magnitude of crystal field splitting depends on strength of the ligands. The strong ligands those appear in spectrochemical series approach closer to the central metal which results in a large crystal field splitting.

• Strong field ligands like CN^—, en, etc. approach closer to the central metal ion, it results in a large crystal field splitting and hence Δ₀ has higher values.

(ii) Oxidation state of the metal : A metal with the higher positive charge is able to draw ligands closer to it than that with the lower one. Thus the metal in higher oxidation state results in larger separation of t_2g and e_g set of orbitals. The trivalent metal ions cause larger crystal field splitting than corresponidng divalent ones.

• For example [Fe(NH₃)₆]³⁺ has higher Δ₀ than [Fe(NH₃)₆]²⁺

Colour of the octahedral complexes:

• A large number of coordination compounds show Wide range of colours due to a’-d transition of electron and this can be explained by crystal field theory (CFT).
• The complex absorbs the light in one visible region (400 nm to 700 nm) and transmits the light in different visible region giving complementary colour.
• Consider an octahedral purple coloured complex of [Ti(H₂O)₆]³⁺ which absorbs green light and transmits purple colour. Similarly [Cu(H₂O)₆] ²⁺ absorbs the light in the red region of radiation spectrum and transmits in the blue region, hence the complex appears blue.
• The absorption of light arises due to d-d transition of electron from lower energy level (t_2g) to higher energy level (e_g) in octahedral -complex.
• The energy required for transition depends upon crystal field splitting energy Δ₀. If Δ₀ = ΔE, then the energy of an absorbed photon (hv) is
• ΔE = hv = hc/λ
• where λ, v and c are wavelength, frequency and velocity of the absorbed light.
• Higher the magnitude of A0 or AE, higher is the frequency or lower is the wavelength of the absorbed radiation.
• Since A0 depends upon nature of metal atom or ion, its oxidation state, nature of ligands and the geometry of the complex, different coordination compounds have different colours.

Octahedral geometry of complexes using crystal field theory :

• In an octahedral complex [MX₆] ^±, the metal atom or ion is placed at the centre of regular octahedron while six ligands occupy the positions at six vertices of the octahedron.
• Among five degenerate d-orbitals, two orbitals namely d_(x² _{– y}²₎ and d_z²; are axial and have maximum electron density along the axes, while remaining three d-orbitals namely d_xy, d_yz and d_zx are planar and have maximum electron density in the planes and in-between the axes.
• Hence, when the ligands approach a metal ion, the orbitals d_(x² _{– y}²₎ and d_z² experience greater repulsion and the orbitals d_xy, d_yz and d_zx experience less repulsion.
• Therefore the energy of d_(x² _{– y}²₎ and d_z² increases while the energy of d_xy, d_yz and d_zx decreases and five d-orbital lose degeneracy and split into two point groups.
• The orbitals d_xy, d_yz and d_zx form t_2g group of lower energy while d_(x² _{– y}²₎ and d_z² form e_g group of higher energy.
• Thus t_2g has three degenerate orbitals while e_g has two degenerate orbitals.
• Experimental calculations show that the energy of t_2g orbitals is lowered by 0.4 Δ₀ or 4D_q and energy of e_g orbital is increased by 0.6Δ₀ or 6D_q. Thus energy difference between t_2g and e_g orbitals is Δ₀ or 10D_q which is crystal field splitting energy.
• CFSE increases with the increasing strength of ligands and oxidation state of central metal ion

Tetrahedral geometry of complexes using crystal field theory :

The tetrahedral structure having the metal atom M at the centre and four ligands occupying the corners is displayed along with in Fig.

• The d_xy, d_yz and d_zx orbitals with their lobes lying in between the axes point toward the ligands. On the other hand, d_(x² _{– y}²₎ and d_z² orbitals lie in between metal-ligand bond axes.
• The d_xy, d_yz and d_zx orbitals experience more repulsion from the ligands compared to that by d_(x² _{– y}²₎ and d_z² orbitals.
• Due to larger such repulsions the d_xy, d_yz and d_zx orbitals are of higher energy while the d_(x² _{– y}²₎ and d_z² orbitals are of relatively lower energy.
• Each electron entering in one of d_xy, d_yz and d_zx orbitals raises the energy by 4D_q whereas that accupying d_(x² _{– y}²₎ and d_z² orbitals lowers it by 6D_q compared to the energy of hypothetical degenerate d orbitals in the ligand field.
• A splitting of d orbitals in tetrahedral crystal fields (assumed to be 10D_q) thus is much less (typically 4/9) compared to that for the octahedral environment.
• The crystal field splitting of d orbitals in a tetrahedral ligand field is compared with the octahedral one in below Fig.

• Thus the pairing of electrons is not favoured in tetrahedral structure. For example, in d⁴ configuration an electron would occupy one of the t_2g orbitals. The low spin tetrahedral complexes thus are not found.
• Typically metal complexes possessing the cetral metal ion with d8 electronic configuration, for example, Ni(CO)₄, favours the tetrahedral structure.

Applications of coordination compounds

In biology :

• Several biologically important natural compounds are metal complexes. They play important role in a number of processes occuring in plants and animals.
• For example, chlorophyll present in plants is a complex of Mg. Haemoglobin present in blood is a complex of iron.

In medicines :

• Pt complex cisplatin is used in the treatment of cancer.
• EDTA is used for treatment of lead poisoning.

To estimate hardness of water :

• Hardness of water is due to the presence of Ca²⁺ and Mg²⁺ ions.
• The ligand EDTA forms stable complexes with Ca²⁺ and Mg²⁺.
• It can, therefore, be used to estimate hardness.

Electroplating :

• This involves deposition of a metal on the other metal. For smooth plating, it is necessary to supply continuously the metal ions in small amounts.
• For this purpose, a solution of a coordination compound is used which dissociates to a very less extent.
• For example, for uniform and thin plating of silver and gold, the complexes K[Ag(CN)₂] and K [Au(CN)₂] are used.