Redox Titrations

Determining the Concentration of an Analyte

As with acid-base titrations, a redox titration (also called an oxidation-reduction titration) can accurately determine the concentration of an unknown analyte by measuring it against a standardized titrant. A common example is the redox titration of a standardized solution of potassium permanganate (KMnO₄) against an analyte containing an unknown concentration of iron (II) ions (Fe²⁺). The balanced reaction in acidic solution is as follows:

$MnO_4^-(aq)+5Fe^{2+}(aq)+8H^+(aq) \rightarrow 5Fe^{3+}(aq)+Mn^{2+}(aq)+4H_2O(l)$

In this case, the use of KMnO₄ as a titrant is particularly useful, because it can act as its own indicator; this is due to the fact that the KMnO₄ solution is bright purple, while the Fe²⁺ solution is colorless. It is therefore possible to see when the titration has reached its endpoint, because the solution will remain slightly purple from the unreacted KMnO₄.

Sample Calculation from Experimental Data

A standardized 4 M solution of KMnO4 is titrated against a 100 mL sample of an unknown analyte containing Fe²⁺. A student conducts the redox titration and reaches the endpoint after adding 25 mL of the titrant. What is the concentration of the analyte?

We know from our balanced equation above that permanganate and iron react in a 1:5 mole ratio. We can therefore perform the following calculation:

Stoichiometric calculation to determine moles of Fe2+

From the balanced equation, Fe²⁺ and KMnO₄ react in a 5:1 mole ratio.

Now that we know the number of moles of iron present in the sample, we can calculate the concentration of the analyte:

$M=\frac {mol}L= \frac {0.5 mol}{0.100 L}=5 M$

Other Types of Redox Titrations

There are various other types of redox titrations that can be very useful. For example, wines can be analyzed for sulfur dioxide using a standardized iodine solution as the titrant. In this case, starch is used as an indicator; a blue starch-iodine complex is formed in the presence of excess iodine, signaling the endpoint.

Another example is the reduction of iodine (I₂) to iodide (I⁻) by thiosulphate (S₂O₃²⁻), again using starch as the indicator. This is essentially the reverse titration of what was just described; here, when all the iodine has been reduced, the blue color disappears. This is called an iodometric titration.

Most often, the reduction of iodine to iodide is the last step in a series of reactions in which the initial reactions are used to convert an unknown amount of the analyte to an equivalent amount of iodine, which can then be titrated. Sometimes halogens (or organic compounds containing halogens) other than iodine are used in the intermediate reactions because they are available in better-measurable standard solutions or they react more readily with the analyte. While these extra steps make an iodometric titration much more involved, they are often worthwhile, because the equivalence point involving the bright blue iodine-starch complex is more precise than various other analytical methods.