Co-ordination compounds or complexes Presented By:- Mrs. Gurvinder Kaur PGT (Chemistry) K V SURA NUSSI

topics Introduction Nomenclature Isomerism Bonding in coordination compounds Colour of complexes Stability of complexes

introduction K4[Fe(CN)6] Central metal ion Ligand Counter ion Co-ordination no. Co-ordination sphere

A Coordination compound consist of a complex ion and necessary counter ions. For example - [Co(NH3)5Cl]Cl2 Complex ion: [Co(NH3)5Cl]2+ Counter ions: 2 Cl-

Complex ion is a species where transition metal ion is surrounded by a certain number of ligands. Transition metal ion : Lewis acid Ligands : Lewis bases e.g. [Co(NH3)6]3+ [Pt(NH3)3Br]+

The molecules or ions coordinating to the metal are the Ligands. They are usually anions or polar molecules having lone pair of electrons to donate. Example :

Monodentate ligands bond using the electron pairs of a single atom. Bidentate ligands bond using the electron pairs of two atoms. Polydentate ligands bond using the electron pairs of many atoms Chelating ligands includes bidentate, Polydentate ligands capable of forming ring with the metal atom or ion.

Polydentate Ligands • Some ligands have two or more donor atoms. • These are called polydentate ligands or chelating agents. • In ethylenediamine, NH2CH2CH2NH2, represented here as en, each N is a donor atom. • Therefore, en is bidentate

The formation of a coordinate complex is a Lewis acid-base reaction Lewis acid: Co3+ Lewis base: Lewis base i.e., the electron pair donor ligand atom is coordinated to a Lewis acid i.e., the electron pair acceptor metal atom or ion. The number of ligand bonds to the central metal atom is termed the coordination number

Nomenclature Cation Anion Naming complex ions: Ligands in alphabetical order Central metal ion Oxidation no. of metal ion in Roman nos.

Examples : K[PtCl3(NH3)] : Potassium amminetrichloridoplatinate(II) K3[Fe(C2O4)3] : Potassium trioxalatoferrate(III) [CoCl2(en)2]Cl : dichloridobis(ethane-1,2-diamine) cobalt(III)chloride [Pt(NH3)4][PtCl4] : tetraammineplatinum(II) tetrachloridoplatinate(II)

BONDING IN CO-ORDINATION COMPOUNDS VALENCE BOND THEORY CRYSTAL FIELD THEORY

Features of valence bond theory : Presence of vacant orbitals Hybridization of suitable number of orbitals Co-ordinate covalent bond between ligand and metal atom Outer or inner orbital complex depending on d-orbitals used in hybridisation

Features of valence bond theory: contd… Valence Bond Theory predicts that in metal complexes bonding arises from overlap of filled ligand orbitals and vacant metal orbitals. Resulting bond is a coordinate covalent bond.

What determines why a metal takes one of these shapes? Complex ions-Three common structural types Octahedral: Most important Tetrahedral Square planar What determines why a metal takes one of these shapes?

Features of crystal field theory : Ligands and metal ion are point charges Electrostatic attraction between ligands and metal ion Splitting of d-orbitals in the presence of ligand fields Nature of ligands affects the extent of splitting

Crystal Field Splitting Splitting of d-orbitals depends on Nature of Ligand Metal Oxidation state Group Position Geometry of the coordination entity

Nature of ligands affects the extent of splitting Stronger the ligand greater is the splitting between d-orbitals High spin Low spin

Metal Oxidation state Greater Oxidation no., (e.g. Fe+3 vs Fe+2) attracts Ligand in closer and results in greater Ligand-Metal interaction and hence greater splitting of d-orbitals

Group Position Further down group, greater Δ d orbitals more diffuse, interact more with ligand

Geometry of the coordination entity Δt = 4/9 Δo

The five d-orbitals in an octahedral field of ligands

Splitting of d-orbitals in octahedral field

Splitting of d-orbital energies by a tetrahedral field of ligands.

Splitting of d-orbital energies by a square planar field of ligands.

Square Planar & Linear Complexes Approach along x-and y-axes Approach along z-axis

ISOMERISM IN COMPLEXES Isomers: same atomic composition, different structures

Ionization isomers [Co(NH3)Br]SO4 & [Co(NH3)SO4]Br [PtCl2(NH3)]Br2 & [PtBr2(NH3)]Cl2

Hydrate isomers Water in outer sphere i.e., water is part of solvent. Water in the inner sphere i.e., water is a ligand in the coordination sphere of the metal.

Linkage isomers Bonding to metal may occur at the S or the N atom Example: Bonding to metal may occur at the S or the N atom Bonding occurs from N atom to metal Bonding occurs from S atom to metal

Linkage Isomers [Co(NH3)5(NO2)]Cl2 Pentaamminenitrito-Ncobalt(III) chloride [Co(NH3)5(ONO)]Cl2 Pentaamminenitrito-Ocobalt(III) chloride

Co-ordination isomers [Co(NH3)6][Cr(CN)6] & [Cr(NH3)6][Co(CN)6] [Pt(NH3)4][PtCl4] & [PtCl(NH3)3][PtCl3(NH3)]

Stereoisomerism

Trans Cis

Geometrical isomerism is not shown by Complexes with co-ordination no.4 having tetrahedral geometry. Square planar complexes of the type MA4,MA3B,MAB3 Octahedral complexes of the type MA6,MA5B.

Example : [Co(NO2)3(NH3)3] Octahedral complexes of the type MA3B3 also exist in two geometrical isomers: fac- and mer-. Example : [Co(NO2)3(NH3)3] NO2 NH3 H3N Co NO2 NH3 H3N Co fac- isomer mer- isomer

fac-isomer mer-isomer

Stereoisomerism Optical isomerism: Have opposite effects on plane-polarized light (non superimposable mirror images)

Enantiomers: non superimposable mirror images A structure is termed chiral if it is not superimposable on its mirror image Mirror image Of structure Structure Two chiral structures: non superimposable mirror images: Enantiomers!

Chirality: the absence of a plane of symmetry Enantiomers are possible A molecule possessing a plane of symmetry is achiral and superimposible on its mirror image.Enantiomers are NOT possible Are the following chiral or achiral structures? Plane of symmetry Achiral (one structure) No plane of symmetry Chiral (two enantiomer)

Which are enantiomers (non-superimposable mirror images) and which are identical (superimposable mirror images)?

Colour in complexes Bonding Theories attempt to account for the colour and magnetic properties of transition metal complexes.

Colour in co-ordination compounds : Due to d-d transition Example: [Ti(H2O)6]3+ Ti3+ : 3d1

Light Absorption If E separating d orbitals lies in visible light range, and d orbitals are not full e- can absorb visible light and jump from a lower orbital to a higher orbital Color perceived is complement of remaining light (VIBGYOR–light absorbed=color perceived)

[V(H2O)6]2+ [V(H2O)6]3+ [Cr(NH3)6]3+ [Cr(NH3)5Cl]2+

Light Absorption Color preceived is on Opposite side from Color absorbed. [Cu (H2O)4]+2 absorbs Orange, appears blue

Magnetic properties of complexes • Crystal Field Splitting can also account for magnetic properties in terms of high-spin and low-spin complexes. • Weak-field ligands lead to high-spin paramagnetic systems. • Strong-field ligands lead to low-spin diamagnetic systems

In terms of MO theory we visualize the coordination as the transfer of electrons from the highest occupied valence orbital (HOMO) of the Lewis base to the lowest unoccupied orbital (LUMO) of the Lewis acid Lewis base Lewis acid Co3+

Lewis acids and bases A Lewis base is a molecule or ion that donates a lone pair of electrons to make a bond A Lewis acid is a molecule of ion that accepts a lone pair of electrons to make a bond Examples : Examples: Electrons in the highest occupied orbital (HO) of a molecule or anion are the best Lewis bases Molecules or ions with a low lying unoccupied orbital (LU) of a molecule or cation are the best Lewis acids

Stability of complexes For some complex ions, the coordinated ligands may be substituted for other ligands Complexes that undergo very rapid substitution of one ligand for another are termed labile Complexes that undergo very slow substitution of one ligand for another are termed inert [Ni(H2O)6]2+ + 6 NH3  [Ni(NH3)6]2+ + 6 H2O (aq)

Stability of complexes contd… M + 4L ML4 KS = [ML4] [M][L]4 The larger the numerical value of KS the more stable is the complex.

The value of KS depends on: charge and size of metal ion charge distribution of metal ion: Class a acceptors i.e.,highly electropositive metals form stable complexes with ligands whose donar atoms are N,O and F. Class b acceptors i.e.,less electropositive and heavier transition metals form stable complexes with ligands whose donar atoms are heavier members of N,O and F groups. basic strength of ligand and its chelate effect

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