1. COORDINATION COMPOUNDS
COMPLEX IONS

2. COORDINATION COMPOUNDS
CoCl3 6NH3
[Co(NH3)6]Cl3
Alfred Werner introduced the 2 types of valences:
Primary valence – oxidation number/charge
Secondary valence – coordination number

3. Metal-Ligand interaction forms a COORDINATE COVALENT BOND
COMPLEXES
Metal – usually transition metal either NEUTRAL or POSITIVELY CHARGED acting as LEWIS ACID
Ligand – usually has at least one pair of unshared valence electrons and acts as LEWIS BASE
Metal-Ligand interaction forms a COORDINATE COVALENT BOND

4. LIGANDS DENTICITY (monodentate, bidentate, polydentate)
CHELATING AGENTS (for bi/polydentate ligands)
FIRST COORDINATION SPHERE as signified by []

5. EXAMPLES [Cu(NH3)2(H2O)Cl]+ [Fe(H2O)2(CN)4]- [Ni(H2NCH2CH2NH2)2]2+
Identify Primary Valence and Secondary Valence

6. NOMENCLATURE

7. RULES Cation Before Anion
Ligand before Central Metal, reverse for formula
Ligand Anions end in –o/-ido
Neutral ligands retain their names
Special Ligand Names
[Co(en)3]Cl3, K2[CoCl4]
Cl-
ethylenediamine
H2O, NH3, CO, NO

8. RULES Greek prefixes (di/bis)
Oxidation number in Roman numeral after the Metal Name
If complex is anion, end in –ate added to Latin name
For multiple ligands, use alphabetical order regardless of prefix
[CoCl4]2-, [Ni(en)2]2+
[Co(NH3)3(NO2)3]

9. EXAMPLE [Ni(CO)4] [Co(H2O)4Cl2]Cl Na3[Ag(S2O3)2] [Fe(en)3](NO3)3
[Cr(NH3)4[FeF6]2

10. SEATWORK [Ag(NH3)2]Cl [Co(NH3)3Cl3] K4[Fe(CN)6] [Ni(CO)4] [Cu(en)2]SO4
[Pt(NH3)][PtCl6]
[CoCl(NH3)4(H2O)]Cl2

11. STRUCTURAL ISOMERS
-compounds of same empirical formula but with different arrangement
4 TYPES
Ionization {[Co(NH3)5(SO4)]Br, [Co(NH3)5Br]SO4}
Hydrate {[Cr(H2O)6]Cl3, [Cr(H2O)5Cl]Cl2H2O}
Linkage {-NO2, -ONO}
Coordination {[Cu(NH3)4][PtCl4], [Pt(NH3)4][CuCl4]}

12. STEREOISOMER -with different spatial arrangement
GEOMETRIC – cis and trans {Co(NH3)4Cl}
OPTICAL – nonsuperimposable mirrors of each other

13. VALENCE BOND THEORY Bond is formed with overlap of two orbitals
For complexes: overlap of ligand orbital (containing an electron pair) and a metal orbital (empty)
Number of Ligands  GEOMETRY
Linear (2), Square Planar (4), Tetrahedral (4), Octahedral (6)

14. EXAMPLES OCTAHEDRAL COMPLEXES [Cr(H2O)6]3+ (outer orbital complex)
[FeF6]3- (inner orbital complex)
SQUARE PLANAR
[Ni(CN)4]2-
TETRAHEDRAL
[CoCl4]2-
H2O and F- with d4 or d6 – inner
D8 – usually outer

15. CRYSTAL FIELD THEORY
Electrostatic attraction (between metal and ligand)
Electrostatic repulsion (between electrons sharing an orbital)
Electrical repulsion is dependent on orientation of d orbitals and the incoming ligands causing splits in energy
crystal field splitting (Δ) – dependent on metal and nature of ligand, affects color and magnetic properties

16. GEOMETRY

17. OCTAHEDRAL COMPLEX
___ ___ ___ ___ ___ eg
Δ(high)
t2g eg
Δ(low)
t2g

18. TETRAHEDRAL COMPLEX
___ ___ ___ ___ ___
t2g
Δ(high) eg

19. SQUARE PLANAR COMPLEX
___ ___ ___
d x2-y2
d xy
d z2
d xz
d yz

20. SPECTROCHEMICAL SERIES
Order of ligands based on ability to produce large Δ (strong field vs weak field) I-<Br-<Cl-<F-<OH-<H2O<NH3<en<NO2-<CN-

21. PARAMAGNETISM and COLOR
Presence of unpaired electrons of the metal
May usually lead to colored complexes
E = hc/λ
Fe3+ in [Fe(H2O)6]3+ vs Ca2+ or Cd2+
[Fe(CN)6]3- vs [FeF6]3-
[CoCl6]3- vs [Co(NH3)6]3+

22. SAMPLE PROBLEM
A compound contains 21.35% Cr, 28.70% N, 6.209% H and 43.68%Cl by mass.
It does not react with HCl.
On reaction with AgNO3, it gives 2 moles of AgCl per mole of compound.
It has an electrical conductivity corresponding to 3mols of ions per mole of compound.
Give the inferences of a, b and c.
Write the formula and give the name of the compound.