Chapter 23: Redox

• It takes electrons from the species being oxidised • It contains the species being reduced

• It gives electrons to the species being reduced • It contains the species being oxidised

• Oxidising/ reducing agents are not just the element that is reduced/ oxidised, but the whole molecule/ ion

• There has to be a change in oxidation number

• By using half equations • By using oxidation numbers

• Write the half-equations for oxidation and reduction, making sure both are balanced • The number of electrons in both equations needs to be the same so the equations can be combined • This may mean that one or both need to be scaled up

• You find which species had a change in oxidation number, and by how much • You scale up the other species (wherever it appears in the reaction- watch out for any that are diatomic) so that the overall change in oxidation number sums to 0

• Start by balancing elements other than hydrogen and oxygen (especially any with a change in oxidation number) • Balance oxygen by adding water (take note of any oxygens already present) • Balance hydrogens by adding protons • Add electrons to balance out charges if necessary (equations also have to be balanced in terms of charge)

• Start by balancing elements other than hydrogen and oxygen (especially any with a change in oxidation number) • Add water molecules to balance oxygens • Add protons to balance hydrogens • Add the same number of protons in hydroxide ions to both sides of the equation • Protons and hydroxide ions react together to form water • Add electrons to balance the equation

• It is usually either water, a proton or a hydroxide ion

• Through redox titrations

• Fe2+/MnO4– • I2/S2O32−

• Standard solution of potassium manganate (VII) is added to the burette. • A pipette is used to add a measured volume of the reducing agent, such as a solution containing iron (II), to the conical flask with an excess of dilute sulfuric acid (as it is needed for the reduction of MnO4-). • No indicator is needed as the reaction is self-indicating; manganate (VII) is purple while manganese (II) is colourless. • The end point is judged by the first permanent pale pink colour, as it shows an excess of manganate (VII) is present (goes from colourless to pale pink). • Burette readings are taken from the top of the meniscus, as potassium manganate (VII) has a deep purple colour that makes it hard to take readings from the bottom of the meniscus. • Repeat until you obtain concordant results.

• Manganate (VII) is reduced to manganese (II), and iron (II) is therefore oxidised to iron (III) • The reactions for both equations can be written, balanced and combined (iron’s would need to be scaled up by a factor of 5) • This tells you they react in a 1:5 ratio • MnO4- + 5Fe2+ + 8H+ -> Mn2+ + 5Fe3+ + 4H2O

• Oxidation: 2S2O3^2- -> S4O6^2- + 2e- • Reduction: I2 + 2e- -> 2I- • Overall: 2S2O3^2- + I2 -> 2I- + S4O6^2-

• Add a standard solution of sodium thiosulfate to the burette. • Prepare a solution of an oxidising agent that you want to analyse, and add it to a conical flask along with an excess of potassium iodide (to supply iodide ions). This will make the solution yellow-brown as iodide ions are oxidised to iodine. • Add the sodium thiosulfate which reduces iodine to iodide, making the yellow-brown colour fade. • When the solution is pale yellow, add starch to make the end point easier to identify. This turns the solution blue-black. • The end point is reached when the blue-black colour disappears.

• Make a standard solution containing chlorate ions • Pour the solution into a conical flask with potassium iodide and anything else needed for the production of iodine • Titrate against thiosulfate ions