Transition metals are elements which form one or more stable ions which have incompletely filled d orbitals. As the only ions of Scandium (3+) and Zinc (2+) form either have empty or full d sub-shells, they are not considered transition metals, although they are still d-block elements.
In general, transition metals:
- show variable oxidation states
- form coloured ions
- form complex ions
- can act as catalysts both as elements and in compounds
The electronic configurations for the d-block elements are largely as expected, with 2 exceptions (Where [Ar] = 1s2 2s2 2p6 3s2 3p6):
A half full or a full 3d sub-shell gives extra stability to an atom. Therefore, for chromium and copper, the 4s sub-shell is only half filled to give this extra stability.
Transition elements lose the 4s electrons before the 3d electrons when forming ions. As these sub-shells are very close together in energy levels, all transition metals are able to form ions with more than one oxidation number - manganese, for example, has ions in oxidation numbers from +1 to +7!
Transition metal ions are small and charged - they have a high charge density. They therefore attract molecules with lots of electrons, ligands, to form complex ions. Ligands are molecules or ions which bond to metal ions by forming dative covalent bonds with an empty d-orbital. A complex ion is a metal atom surrounded by ligands. The coordination number of a complex ion is the number of dative covalent bonds formed from ligand(s); this is not the same as the number of ligands because some molecules can form more than 1 dative covalent bond to a metal ion. These are called bidentate (e.g. ethane-1,2-diamine or ethanedioate ions) or multidentate (e.g. EDTA) ligands, rather than unidentate (e.g. ammonia, water, chloride) ligands.
Shapes of Complex Ions
With small ligands, such as water, transition metal ions form octahedral complexes. This is because there are 6 ligands around a central metal ion which spread out as far as possible e.g. [Cu(H2O)6]2+, [Cr(NH3)6]3+. With larger ligands, such as chloride ions, this is not enough room for 6 ligands to fit around the metal ion, so a tetrahedral complex with 4 ligands is formed e.g. [CrCl4]-. There are also times when square planar complexes e.g. [Pt(NH3)2Cl2]/cisplatin, along with linear complexes e.g. [Ag(NH3)2]+/Tollens' reagent, or [CuCl2]- are formed. All the complexes here follow the VSEPR theory covered in AS chemistry.
The 5 d orbitals point in different directions. For an isolated ion, all of these orbitals are at the same energy level. When ligands form dative covalent bonds to the metal ion, electrons in some of the orbitals will experience a greater repulsion than electrons in other orbitals. The 3d sub-shell is split, with one set of orbitals at a higher energy level and one at a lower energy level. The energy difference between these is the energy of a photon given out when electrons move to the lower energy orbitals. The energy difference and photon frequency is linked by:
Electrons can only move between these energy levels when the d sub shell is neither full nor empty - this explains why scandium and zinc do not have coloured complexes. The energy difference depends on the oxidation state of the metal ion, the coordination number and number of ligands around the metal ion.
Transition metal complexes are also able to absorb light when electrons move into the higher energy orbitals. This is useful spectrometry (AQA only)
Transition metal ions can be identified by colour. The table below shows the aqua complexes of the transition metals for the common oxidation states:
|+3 colourless||+3 violet||+3 blue||+3 ruby||+2 pale pink||+2 pale green||+2 pink||+2 green||+1 colourless||+2 colourless|
|+5 yellow||+6 yellow or orange1||+7 purple||+3 red-brown|
1Chromate (CrO4-) = yellow, dichromate (Cr2O72-) = orange
Reactions of Complexes (Edexcel)
The following table shows the results of adding aqueous sodium hydroxide and ammonia to some transition metal ions:
|Ion||Some NaOH or NH3||Excess NaOH||Excess NH3|
|Cr(III)||Green or violet precipitate||Deep green solution||No change|
|Mn(II)||Buff precipitate which darkens in air as MnO2 forms||No change||No change|
|Fe(II)||Pale green precipitate which turns brown at surface due to oxidation to Fe(III)||No change||No change|
|Fe(III)||Red-brown precipitate||No change||No change|
|Ni(II)||Pale green precipitate||No change||Pale lavender-blue solution|
|Cu(II)||Pale blue precipitate||No change||Deep blue solution|
|Zn(II)||White precipitate||Colourless solution||Colourless solution|
The precipitates formed upon addition of some NaOH or NH3 all involve the formation of the hydroxide of the metal when hydrogen ions are pulled off the hexaaqua complexes of the ions in solution by the hydroxide ions from NaOH or NH4OH, e.g.:
The deep green solution of chromium in excess NaOH is due to:
The pale solution of nickel in excess NH3 is due to:
The blue solution of copper in excess NH3 is due to:
The colourless solutions of zinc are due to:
Use as Catalysts
Transition metals are used in industry as catalysts due to their variable oxidation states from the ability of the partially filled 3d orbitals to gain or lose electrons.
Heterogeneous Catalysts (AQA)
Heterogeneous catalysts are ones which exist in a different phase from the reactants. Examples include:
- Pt, Pd, Rh in catalytic converters. These solid catalysts are supported on a ceramic honeycomb to maximise the surface area and minimise cost.
- V2O5 in the Contact Process:
Vanadium changes between the +5 and +4 oxidation states.
- Cr2O3 to manufacture methanol from carbon monoxide and hydrogen.
- Fe in the Haber Process.
Although catalysts are not used up in the reactions they catalyse, they may get poisoned by impurities in the reactants and lose their efficiency. This means catalysts need to be replaced over time, which may get expensive as some of these metals are rare - e.g. lead poisons catalytic converters, and sulfur poisons iron in the Haber Process.
Homogeneous Catalysts and Autocatalysis(AQA)
When a catalyst and the reactants are in the same phase, the catalyst is called a homogeneous catalyst. The catalyst reacts and is regenerated in the reaction:
- Fe2+ catalyses the reaction between I- and S2O82-:
The overall reaction involves 2 negatively charged ions, which repel each other. The reactions with the iron catalyst involve oppositely charged ions, which make it easier for the reaction to take place.
Autocatalysis is when a reaction gives a product which catalyses the reaction itself:
- Mn2+ catalyses the reaction between C2O42- and MnO42-:
The purple colour of manganate(VII) disappears slowly at first, but this increases in speed as more Mn2+ is produced.
Applications of Complexes (AQA)
- Variable oxidation states are important in catalysts.
- Fe(II) in haemoglobin enables oxygen to be transported in the blood. CO is toxic as it binds to the iron very strongly, stopping it from being able to bind to oxygen to carry it around.
- Pt(II) in cisplatin is used as an anticancer drug. Benefits: treats cancers by triggering apoptosis. Risks: major side effects include kidney and nerve damage, hearing and hair loss.
- [Ag(NH3)2]+ is in Tollens' reagent to distinguish between aldehydes and ketones.