UNIT V SEC.V-7 REDOX TITRATION
Redox titrations are used to determine the exact
concentration of a substance that undergoes
oxidation or reduction. Just like an Acid/Base titration
a colour change is used to determine the endpoint of
the redox reaction.
Some species undergo significant colour change when
being oxidized or reduced. Chemists make use of this
colour change to determine the endpoint and to
calculate the concentration of an unknown species.
A common Oxidizing Agent is
(permanganate ion)
MnO4-
This ion is usually made by dissolving the salt KMnO4 in
water. The MnO4-(aq) is a very intense purple colour, but
when it is reduced to the Mn2+ ion it becomes colourless.
MnO4- + 8H+ + 5e-  Mn2+ + 4 H2O
Purple colour
clear colour
This reduction has to be done under acidic conditions
(note the presence of H+) and the drastic colour
change from clear to purple indicated the end point.
Let’s try to see what happens when we use this
oxidizing agent to oxidize Fe2+ ion to Fe3+ ion.
Part 1: Preliminary Investigation of KMnO4 as an Oxidizing agent.
Let’s see what colour the ions are under acidic, basic, and neutral
conditions:
Data Table:
Type of Solution
Colour
Ion or Molecule Present
Acidic (3 M H2SO4)
clear
Mn2+
Neutral (water)
brown
MnO2
Basic (6 M NaOH)
purple
MnO4-
Here is the two half-cell reactions:
2 X MnO4- + 8H+ + 5e-  Mn2+ + 4 H2O
5X
SO32- + H2O  SO42- + 2 H+ + 2 e-
2 MnO4- + 5 SO32- + 6H+ 
2 Mn2+ + 5 SO42- + 3 H2O
Part 2: Determining the concentration of a Solution of Fe2+
The permanganate ion will be titrated against a solution of
FeSO3. The end point is reached when the colourless solution of
FeSO3 remains a slight purple colour indicating that all of the Fe2+
has been reduced to Fe3+ ions.
Here is the two half-cell reactions:
MnO4- + 8H+ + 5e-  Mn2+ + 4 H2O
Fe2+  Fe3+ + 1 e-
1. Write out the balanced Redox reaction for these two
half-cell reactions for acidic conditions.
MnO4- + + 5 Fe2+ + 8H+  Mn2+ + 5 Fe3+ + 4 H2O
2. A 0.200 M KMnO4 solution will then be titrated
against an unknown concentration of acidified FeSO3
solution. 10.0mL of FeSO3 solution will be placed into a
250 mL Erlenmeyer flask and 4.0 mL of 3 M H2SO4 is
then added to this.
3. The endpoint will occur when a permanent slight
purple colour appears in the Erlenmeyer flask.
Data Table: Determining the Concentration of the Solution of Fe2+;
Volume of KMnO4 to react with 10.0 mL of Fe2+ solution
[KMnO4] = ______
Trial #1
Initial Volume of
KMnO4 (mL)
0.00
Final Volume of
KMnO4 (mL)
Volume of KMnO4
required (mL)
Average Volume
(mL)
Trial #2
Trial #3 (If
necessary)
Analysis:
1. Calculate the number of moles of the KMnO4 solution was used
at the end point.
n = C xV
2. How many moles of Fe2+ ions can the number of moles of
KMnO4 solution oxidize? (Hint: Use your balanced REDOX
reaction).
3. What
was the concentration of the Fe2+ solution?
1) Do the Follow up Questions: #1 & #2
2) For more practice go to your workbook p. 213 – 214
and do exercises #26, 28, 30 and 32
3) Quiz next class on Balancing REDOX reactions and
REDOX Titrations