Redox

Redox chemistry describes reactions in which one component is oxidised (loses electrons) and another component is reduced (gains electrons).

There are two well-known, classroom-based redox titrations:

 1. Determining the concentration of Fe²⁺ ions in solution

The underlying chemistry of the redox reaction is shown below.

 5Fe²⁺_(aq) + MnO₄^-_(aq) + 8H⁺_(aq) → 5Fe³⁺_(aq) + Mn²⁺_(aq) + 4H₂O_(l)

For many students, balancing the equation will be the first challenge followed by the multistep calculations to determine the concentration of the Fe²⁺ ions. Mistakes are likely to be made when considering the 5:1 ratio of Fe(II) ions to the manganate ion. It’s also important to draw students’ attention to the part that adding sulfuric acid (H⁺ ions) plays in the overall redox reaction.

2. Determining the concentration of sodium thiosulfate solution

The underlying chemistry of the redox reaction is shown below and takes place in two parts. A redox reaction takes place in the conical flask prior to the titration.

10I^-_(aq) + 2MnO₄^-_(aq) + 16H⁺_(aq) → 5I_2(aq) + 2Mn²⁺_(aq) + 8H₂O_(l)

 A second redox reaction takes place on addition of sodium thiosulfate…

 I_2(aq) + 2S₂O₃^2-_(aq) → S₄O₆^2-_(aq) + 2I^-_(aq)

 …although it’s more accurate to identify that it is likely a triiodide ion is oxidising the sulfur in the thiosulfate ion to the tetrathionate ion.

I^-_(aq) + I_2(aq) → I₃^-_(aq)

I₃^-_(aq) + 2S₂O₃^2-_(aq) → S₄O₆^2-_(aq) + 3I^-_(aq)

The starch indicator is only added near the endpoint of the titration as the starch-triiodide complex can stabilise over time in a high concentration of I₂ and becomes difficult to reduce.

Questions you could ask your students for both redox reactions:

• Which species is being reduced?
• Which species is being oxidised?
• Why don’t we directly titrate iodine solution with sodium thiosulfate?
• Why is no indicator used in the iron tablet titration?
• Why is starch used as an endpoint indicator in the thiosulfate titration?
• What is the oxidation number of manganese in the permanganate ion?
• What is the oxidation number of sulfur in the thiosulfate ion and in the tetrathionate ion?

For more videos on redox titrations, please see Practical videos: titrations.

The following video from Pdst chemistry provides detailed demonstrations of the halogens as oxidising reagents, as well as metal displacement reactions:

• Chlorine with bromide ions
• Chlorine with iodide ions (1:33)
• Bromine with iodide ions (2:13)
• Halogen reactions with sulfite ions (3:58)
• Halogen reactions with Fe²⁺ ions (6:52)
• Displacement reactions of Zn with Cu^2+. (8:39)
• Displacement reaction of Mg with Cu²⁺ (9:52)

In all of these experiments, invite your students to write balanced redox equations. Ensure they grasp what is meant by the terms ‘oxidise’ and ‘reduce’ and which species is acting as the oxidising/reducing agent. For example, here are two different ways to describe the reaction of chlorine with bromide ions:

• Chlorine oxidises the bromide ions to bromine.
• The bromide ion reduces the chlorine to chloride ions.

Finally, ask your students what they would expect to happen if they added bromine to a solution containing chloride ions. Why?