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COORDINATION COMPOUNDS

COORDINATION COMPOUNDS. COMPLEX IONS. COORDINATION COMPOUNDS. CoCl 3  6NH 3 [Co(NH 3 ) 6 ]Cl 3 Alfred Werner introduced the 2 types of valences: Primary valence – oxidation number/charge Secondary valence – coordination number. COMPLEXES.

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COORDINATION COMPOUNDS

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  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. 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/polydentateligands) • 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 – cisand 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-

  15. CRYSTAL FIELD THEORY • Electrostatic attraction (between metal and ligand) • Electrostatic repulsion (between electrons sharing an orbital) • Electrical repulsion is dependent on orientation of dorbitals 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 ___ ___ ___ ___ ___ dx2-y2 dxy dz2 dxz dyz

  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.

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