Inorganic Chemistry and Periodicity
This domain tests the descriptive chemistry that distinguishes groups and periods: s-block trends, halogen behaviour, the characteristic properties of transition metals, and the observations and equations you are expected to recall rather than derive.
Periodic trends
- Group 2 reactivity with water increases down the group as ionisation energy falls. Hydroxide solubility increases down the group while sulfate solubility decreases, which is why BaSO4 is used as a precipitate.
- Thermal stability of Group 2 carbonates and nitrates decreases down the group because the larger cation distorts the anion less.
- Halogen oxidising power decreases down the group; halide reducing power increases. Hydrogen halides become less thermally stable and their aqueous solutions more acidic down the group.
- Disproportionation, for example Cl2 + 2NaOH → NaCl + NaOCl + H2O.
Transition metals
- Defined by a partially filled d subshell in at least one common oxidation state; zinc (d10) is excluded.
- Characteristic properties: variable oxidation states (Mn from +2 to +7), catalytic activity, coloured complex ions and paramagnetism.
- Ligands donate a lone pair to form coordinate bonds. Coordination numbers are commonly 6 (octahedral) and 4 (tetrahedral or square planar).
- Colour arises from d–d transitions split by the ligand field, so d0 and d10 ions such as Sc3+ and Zn2+ are colourless.
Worked example
In [Fe(CN)6]3−, six cyanide ligands at −1 give −6; the overall charge is −3, so iron is +3. Fe3+ is d5. Cyanide is a strong-field ligand, so the configuration is low-spin t2g5 with one unpaired electron - the ion is paramagnetic and coloured.
Traps that cost marks
- Reversing the hydroxide and sulfate solubility trends down Group 2.
- Calling all d-block elements transition metals, including Zn and Sc.
- Assuming every transition-metal ion is coloured.
- Quoting a qualitative observation without a balanced equation.
- Forgetting that halide reducing power increases, not decreases, down the group.