COORDINATION COMPOUNDS COMPLEX IONS

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

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

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

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

NOMENCLATURE

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

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]

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

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

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]}

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

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)

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

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

GEOMETRY

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

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

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

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-

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+

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.