Cobalt(III) Complexes and Werner's Theory

Cobalt(III) Chloride-Ammonia Complexes

• There are four stable cobalt(III) chloride-ammonia complexes with different colours and conductivities.
• The colour and conductivity of the complexes are determined by the number of chloride ions outside the coordination sphere.
• Yellow [Co(NH3)6]3+ 3Cl- is a 1:3 electrolyte, meaning it dissociates into four ions in solution.
• Purple [CoCl(NH3)5]2+ 2Cl- is a 1:2 electrolyte, meaning it dissociates into three ions in solution.
• Green [CoCl2(NH3)4] + Cl- is a 1:1 electrolyte, meaning it dissociates into two ions in solution.
• Violet [CoCl2(NH3)4] + Cl- is a 1:1 electrolyte, meaning it dissociates into two ions in solution.
• The green and violet complexes have the same empirical formula, CoCl3.4NH3, but different properties, making them isomers.

Werner's Theory of Coordination Compounds

• Werner's theory explains the bonding and structure of coordination compounds.
• It describes two types of valences: primary and secondary.
• Primary valences are ionizable and satisfied by negative ions, while secondary valences are non-ionizable and satisfied by neutral molecules or negative ions.
• Secondary valence equals the coordination number and is fixed for a specific metal.
• Ions/groups bound by secondary linkages to the metal have specific spatial arrangements called coordination polyhedra.
• Coordination entities or complexes are enclosed in square brackets, while ions outside the brackets are counter ions.
• Common coordination polyhedra in transition metal complexes include octahedral, tetrahedral, and square planar.

Magnetic Properties of Coordination Compounds

• The magnetic properties of coordination compounds are determined by the number of unpaired electrons in the metal ion.
• Coordination compounds with up to three d electrons have two vacant d orbitals for octahedral hybridisation, resulting in similar magnetic behaviour to their free metal ions.
• When more than three d electrons are present, pairing of d electrons is required for octahedral hybridisation, influencing magnetic behaviour.
• Magnetic data generally agrees with maximum spin pairing, especially for d6 complexes.
• Exceptions exist for d4 and d5 complexes. For example, [Mn(CN)6]3- has two unpaired electrons while [MnCl6]3- has four unpaired electrons.
• Werner's theory has limitations, failing to predict tetrahedral and square planar structures for four-coordinate complexes and to distinguish between weak and strong ligands.

Crystal Field Theory

• Crystal Field Theory (CFT) explains bonding and structure of coordination compounds using an electrostatic model.
• CFT considers the metal-ligand bond as purely ionic, arising from electrostatic interactions between the metal ion and ligands.
• Ligands are treated as point charges for anions or point dipoles for neutral molecules.
• In an isolated metal atom/ion, all five d orbitals are degenerate (have the same energy).
• This degeneracy is lifted when a non-spherical field of negative charges from ligands surrounds the metal atom/ion.
• The splitting pattern of d orbitals depends on the crystal field.
• In octahedral coordination entities, ligands cause repulsion with metal d orbitals, leading to energy splitting.
• d x2-y2 and d z2 orbitals, pointing towards ligands, experience greater repulsion and are raised in energy.
• dxy, dyz, and dxz orbitals, directed between axes, are lowered in energy.

Colour of Coordination Compounds

• The colour of coordination compounds is explained by CFT and relates to the absorption of specific wavelengths of light.
• The colour observed is the complementary colour of the absorbed light.
• For example, [Ti(H2O)6]3+ appears violet because it absorbs blue-green light.
• Light absorption excites an electron from a lower energy level (t2g) to a higher energy level (eg).

Metal Carbonyls

• Homoleptic carbonyls are compounds containing only carbonyl ligands and form readily with most transition metals.

• They exhibit well-defined structures such as tetrahedral, trigonal bipyramidal, and octahedral.

• Examples include:

  • Tetracarbonylnickel(0) - tetrahedral
  • Pentacarbonyliron(0) - trigonal bipyramidal
  • Hexacarbonyl chromium(0) - octahedral
  • Decacarbonyldimanganese(0) - two square pyramidal units joined by a Mn-Mn bond
  • Octacarbonyldicobalt(0) - Co-Co bond bridged by two CO groups

• The metal-carbon bond in metal carbonyls has both σ and π character.

• The σ bond is formed by donation of the lone pair from the carbonyl carbon to a vacant metal orbital.