A transition metal is a metal that can form one or more stable ions with an incomplete d sub-level. This allows for the formation of coloured ions. In addiiton, transition metals have variable oxidation states, which is useful for catalysis. Transition metals also form complexes (see below).
A complex is a central metal atom or ion surrounded by a ligand. A ligand is a molecule or ion that forms a co-ordinate bond with a transition metal by donating a pair of electrons. Different complexes have different shapes depending on the co-ordination number as well as the ligand size. Ligands could be monodente, bidentate or multidentate.
The binding of ligands splits d orbitals apart and d electrons move from the ground state to an excited site when light is absorbed. Transition metal colours arise due to some wavelengths of light being absorbed and the rest reflected. The energy gap between the 2 orbitals can be calculated as E = hf where h is the Planck's constant and f is the frequency.
Ligand substitution is where one ligand within a metal complex is swapped to another ligand. Substitution of similarly sized ligands resolves in no change in the coordination number but substitution of different sized ligands resolves in a change in the coordination number. The chelate effect states that multidentate ligands always form much more stable complexes than monodente ligands due to the increase in entropy.
Vanadium contains ions with different possible oxidation numbers and different species are formed by the reduction of vanadate ions by zinc in acidic solution. Different vanadium species contain different visible colours. The redox potential for a transition metal ion changing from a higher to a lower oxidation state is influenced by pH and by the ligand.
Besides acid-base neutralisation reactions, titrations can also be carried out by transition metal redox reactions. 2 examples of redox titrations include reacting Fe2+ with manganate ions as well as C2O42- with manganate ions.
There are 2 types of catlysts - homogenous catlysts (same phase as reactants) and hetrogeneous catalysts (different phase as reactants). Homogenous catalysts often involve forming an intermediate species to lower the activation energy by having 2 oppositely charged ions reacting together.