Redox titrations similar to acid- base titration.
If don’t know which way round colour changes, look at table to see what’s added.
Manganate (VII) Titrations: Potassium manganate (VII) KMnO4 under acidic conditions.
Potassium manganate can be replaced by other oxidising agents such as acidified dichromate
(VI) H+/Cr2O7
2
.
– MnO4- ions reduced to Mn2+
.
– Other chemical used must oxidise and a reducing agent that reduces MnO4- to Mn2+. Such as
can analyse iron ions Fe2+ or ethanedioic acid (COOH)2.
– No indicator, reaction self- indicating.
– Potassium manganate deep purple. So always read from top of meniscus as can’t see
bottom.
Method:
– Standard solution of potassium manganate added to burette.
– Using pipette, measure 25cm3
of a solution of oxidising agent from 250cm3
volumetric flask
to a conical flask.
– Also add excess dilute sulfuric acid to provide H+ ions for the reduction of MnO4-.
– Manganate solution reacts and is decolourised as being added.
– Goes colourless to pink at end point.
– Repeat titration until obtain concordant tires- two tires that agree +/- 0.1
Analysing Percentage Purity of Iron (ll) Compound:
– Work out mass of impure sample- what I made up in volumetric flask with H2O.
– Work out mean titre- the volume of MnO4- reacted. Only concordant titre used in mean.
– Calculate moles of MnO4- reacted using CV/ 1000.
– Calculate moles of Fe2+ reacted using molar ratio- from overall equation. Scale up this
number to match moles in 250cm3
by multiplying by 10. Moles of Fe2+ = moles of FeSO4.7H2O
= pure.
– Find Mr of FeSO4.7H2O. Find mass of FeSO4.7H2O using Mr x moles.
– Divide this number by mass of impure sample found at the start to find out % purity of
FeSO4.7H2O in impure sample. x 100.
Can also work out x in (COOH)2.xH2O using titration like above. Again moles of (COOH)2.xH2O =
moles of (COOH)2. Same calculation above however in last step would work out molar mass of
H2O needed, which would be the x. Learn all four below.
