• Redox titration involves an?oxidising agent?being titrated against a?reducing agent
• Electrons are transferred from one species to another
• In acid-base titrations indicators are used to show the endpoint of a reaction; however?redox titrations?using transition metal ions naturally change colour when changing oxidation state, so indicators are not always necessary
• They are said to be 'self-indicating'
• Two common transition metal ion redox titrations are?manganate(VII)?and?dichromate(VI)?titrations

Manganate(VII) Titrations

• Potassium manganate(VII)?is an oxidising agent and is a deep purple colour
• In acidic solutions it is reduced to the almost colourless?manganese(II)?ion (the ion is actually pink, but in low concentrations it is effectively colourless)
• The reduction equation for the?manganate(VII)?ion is

MnO4-?(aq) + 8H+?(aq) ?+ 5e-? ?→? ?Mn2+?(aq) ?+ 4H2O (aq)

purple? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? colourless

• The potassium ion is a spectator ion so can be left out of the equations
• By convention the?potassium manganate(VII)?solution is placed in the burette, so that as it reacts with the reducing agent the solution becomes colourless
• At the endpoint the?manganate(VII)?ion becomes in excess so the first appearance of a permanent colour change marks the endpoint
• The colour seen is pink; if the colour is purple then you have overshot the endpoint and there is too much?manganate(VII)?in excess
• Analysis of iron in iron(II)sulfate tablets is typical manganate(VII) titration

Answer

Step 1: Write the balanced equation for the reaction

oxidation: Fe2+?(aq) ?→? ?Fe3+?(aq) ?+?e-

reduction: MnO4-?(aq) + 8H+?(aq) ?+ 5e-? ?→? ?Mn2+?(aq) ?+ 4H2O (l)

overall:?MnO4-?(aq) + 8H+?(aq) ?+ 5Fe2+?(aq)? ?→? ?Mn2+?(aq) ?+ 4H2O (l) + 5Fe3+?(aq)

Step 2: Determine the amount of MnO4-?used in the titration

moles of MnO4-??= 0.0265 dm3??x? 0.100 mol dm-3?=?0.00265 mol

Step 3:?Determine the amount of iron in the reaction

From the equation for the reaction we know the reacting ratio??MnO4-?:?Fe2+?= 1: 5

∴ moles of?Fe2+?=?0.00265 mol MnO4-?x 5 =?0.01325 mol

Step 4:?Convert moles into mass of iron

Mass of iron =?0.01325 mol x 55.85 gmol-1?=?0.740 g

Step 5:?find the percentage of iron in the tablet

∴ % Fe in the tablet = (0.740/ 2.25)? x 100 =?32.9%

Dichromate(VI) Titrations

• Potassium?dichromate(VI),?K2Cr2O7, is another oxidising agent used in redox titrations
• The oxidation state of the chromium changes from +6 to +3 in the reaction:

Cr2O7-?(aq) + 14H+?(aq) ?+ 6e-? ?→? ?2Cr3+?(aq) ?+ 7H2O (l)

orange? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? ? green

• From the half equation you can see that the reaction needs hydrogen ions so the solution must be acidified with excess dilute sulfuric acid
• The colour change is hard to see as the?dichromate(VI)?ion is pale orange and the?chromium(III)?ion is pale green
• To enhance the endpoint an indicator called diphenylaminesulfonate is used which turns from colourless to purple at the endpoint
• To obtain the balanced equation in the analysis of iron(II) first work out the redox changes in the half equations, then balance the number of electrons transferred:

The steps in producing a balanced redox equation between iron(II) and dichromate(VI)