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

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

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

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

5. 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]+

6. 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 :

7. 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.

8. 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

10. 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

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

12. 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)

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

14. 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

15. 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.

16. 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?

20. 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

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

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

24. 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

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

26. Geometry of the coordination entity
Δt = 4/9 Δo

27. The five d-orbitals in an octahedral field of ligands

28. Splitting of d-orbitals in octahedral field

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

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

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

34. ISOMERISM IN COMPLEXES
Isomers: same atomic composition, different structures

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

36. 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.

37. 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

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

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

40. Stereoisomerism

41. Trans
Cis

42. 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.

43. 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

44. fac-isomer
mer-isomer

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

46. 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!

47. 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)

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

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

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

51. 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)

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

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

57. 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

58. 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+

59. 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

60. 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] NH3  [Ni(NH3)6] H2O (aq)

61. 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.

62. 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

63. Thank You !