manganate(VII) titrations;
non-structured calculations based on these and any other types of titration

Redox titrations

the term coordination number, the shapes and bond angles of complexes with coordination numbers 4 (square planar and tetrahedral) and 6 (octahedral)

balancing half-equations and full equations for redox processes that also include acid-base reactions by using oxidation states or other methods

e.g. \(MnO_4^- + 5e^- + 8H^+ → Mn^{2+} + 4H_2O\) and
\(MnO_4^- + 5Fe^{2+} + 8H^+ → Mn^{2+} + 5Fe^{3+} + 4H_2O\)
This supplements ES(f).

Redox titrations

simple electrochemical cells:
(i) involving metal ion/metal half-cells
(ii) involving half-cells based on different oxidation states of the same element in aqueous solution with a platinum or other inert electrode, acidified if necessary
(iii) techniques and procedures to set up and use electrochemical cells

Introduction to redox equilibria Standard hydrogen electrode Standard electrode potentials Combining half-cells

the action of an electrochemical cell in terms of half-equations and external electron flow and the ion flow in the salt bridge

Combining half-cells

the term standard electrode potential and its measurement using a hydrogen electrode;
use of standard electrode potentials to:

(i) calculate E_cell
(ii) predict the feasibility of redox reactions, and the reasons why a reaction may not occur
(iii) explain rusting, and its prevention, in terms of electrochemical processes

Details of the set-up of the hydrogen electrode are not required, just the equation for the reaction.
Learners should know the standard conditions.

Standard hydrogen electrode Standard electrode potentials Combining half-cells

transition metals as d-block elements forming one or more stable ions which have incompletely filled d-orbitals;
the common oxidation states of iron (+2 and +3) and copper (+1 and +2) and the colours of their aqueous ions, if any

Learners should also be able to use given data about transition metals and their compounds.

electronic configurations, using sub-shells and atomic orbitals, for ions of the first row of the d-block elements;
the existence of variable oxidation states, in terms of the stability of d-orbital electron arrangements

The electron configurations of Cu and Cr may be required.
See also EL(f).

the terms ligand, complex/complex ion and ligand substitution

Learners should know the formulae of the following examples of complex ions from the chemistry of:
iron: [Fe(H₂O)₆]²⁺, [Fe(H₂O)₆]³⁺;
and copper: [Cu(H₂O)₆]²⁺, [Cu(NH₃)₄]²⁺, [CuCl₄]^2-.
They should be able to write similar formulae for other complexes, given suitable information.

the formation of complexes in terms of coordinate (dative) bonding between ligand and central metal ion;
ligand substitution equations;
the terms bidentate and polydentate as applied to ligands

Learners should know the structure of ethanedioate and how it acts as a bidentate ligand.
Formulae of other multidentate ligands will be given.

the colour changes in, and ionic equations for, the reactions of: Fe²⁺(aq), Fe³⁺(aq) and Cu²⁺(aq) ions with sodium hydroxide solution and ammonia solution

[Iron(II) hydroxide and iron(III) hydroxide do not form complexes with ammonia.]

the catalytic activity of transition metals and their compounds

Homogeneous catalysis in terms of variable oxidation states.
Heterogeneous catalysis in terms of the ability of transition metals to use 3d and 4s electrons of the atoms on the catalyst surface to form weak bonds to reactants.

(i) the ions of transition metals in solution are often coloured;
(ii) the origins of colour in transition metal complexes in terms of the splitting of the d-orbitals by the ligands and transitions between the resulting electronic energy levels

Details of how the d-electrons split in a particular complex are not required.

techniques and procedures to measure concentrations of solutions using a colorimeter or visible spectrophotometer