Q.Which method can be used to find out strength of reductant/oxidant in a solution? Explain with an example.

The key idea to determine the strength of a reductant or oxidant in a solution is Redox Titration.

Redox titration is a volumetric analysis method where a solution of known concentration (the titrant, either an oxidant or a reductant) is used to determine the unknown concentration of another solution (the analyte, the opposite species). The process involves:

1. A precisely measured volume of the analyte is taken.
2. The titrant is added gradually from a burette until the equivalence point is reached, where the redox reaction between the titrant and analyte is complete.
3. The equivalence point is detected by a sharp colour change, either from a redox indicator or by the titrant itself if it is coloured (self-indication).
4. Using the volume of titrant consumed, its known concentration, and the stoichiometry of the balanced redox reaction, the concentration (and thus the strength) of the unknown analyte can be calculated.

Example:

To find the strength of a ferrous salt solution (a reductant), it can be titrated against a standard solution of potassium permanganate ($\text{KMnO}_4$, a strong oxidant) in an acidic medium.

The balanced ionic reaction is:

$5\text{Fe}^{2+} + \text{MnO}_4^- + 8\text{H}^+ \rightarrow 5\text{Fe}^{3+} + \text{Mn}^{2+} + 4\text{H}_2\text{O}$ …

Concept and Intuition

Reductants (reducing agents) are chemical species that donate electrons and get oxidised themselves, causing another species to be reduced. Conversely, oxidants (oxidising agents) are species that accept electrons and get reduced themselves, causing another species to be oxidised. The "strength" of a reductant or oxidant refers to its inherent tendency to lose or gain electrons, respectively. In a practical solution, this strength is also directly related to the concentration of the active species available to participate in redox reactions.

To quantify this available oxidising or reducing power in a solution, we need a method that can accurately measure the amount of the active species present. Redox titration is the primary analytical technique used for this purpose. It allows us to determine the unknown concentration of a reductant or oxidant by reacting it completely with a precisely known amount of a complementary oxidant or reductant.

Method: Redox Titration

Redox titration is a type of volumetric analysis based on a redox reaction between the analyte (substance of unknown concentration) and the titrant (substance of known concentration).

1. Principle: The core idea is to react a solution containing the unknown reductant/oxidant (analyte) with a solution of a known concentration of an oxidant/reductant (titrant) until the reaction is complete. By measuring the volume of the titrant required to reach this completion point, and knowing the stoichiometry of the reaction, the concentration of the analyte can be calculated.

2. Key Components:

   • Titrant: The solution of precisely known concentration (a standard solution) that is added from a burette.
   • Analyte: The solution of unknown concentration, typically placed in a conical flask.
   • Indicator: A substance that undergoes a distinct colour change at or very near the equivalence point, signaling the completion of the reaction (the end point). In some redox titrations, one of the reactants itself can serve as a self-indicator.
   • Equivalence Point: The theoretical point in the titration where the amount of titrant added is chemically equivalent to the amount of analyte initially present, according to the stoichiometry of the balanced redox reaction.
   • End Point: The experimental point where the indicator changes colour. Ideally, the end point should be very close to the equivalence point.

3. Procedure:

   • A precise, known volume of the analyte solution (of unknown concentration) is measured using a pipette and transferred into a conical flask.
   • The burette is filled with the titrant (of known concentration).
   • The titrant is slowly added dropwise from the burette to the analyte in the conical flask, with continuous swirling, until the indicator changes colour permanently. This marks the end point.
   • The volume of titrant consumed is recorded from the burette.
   • The titration is usually repeated several times to ensure accuracy and precision, and an average volume is used for calculations.

4. Calculation: Once the volume of titrant consumed is known, along with its concentration and the volume of the analyte, the unknown concentration of the analyte can be calculated using the stoichiometry of the balanced redox reaction.

   For a general redox reaction where $n_1$ moles of reactant 1 react with $n_2$ moles of reactant 2:

   $\frac{M_1 V_1}{n_1} = \frac{M_2 V_2}{n_2}$

   Where $M_1, M_2$ are the molarities, $V_1, V_2$ are the volumes, and $n_1, n_2$ are the stoichiometric coefficients from the balanced equation.

   Alternatively, using normality (where $N = M \times \text{n-factor}$):

   $N_1 V_1 = N_2 V_2$

Example: Determining the Strength of Ferrous Ammonium Sulphate Solution

Let's consider the determination of the strength (concentration) of a ferrous ammonium sulphate (Mohr's salt, $\text{FeSO}_4 \cdot (\text{NH}_4)_2\text{SO}_4 \cdot 6\text{H}_2\text{O}$) solution using a standard solution of potassium permanganate ($\text{KMnO}_4$). Here, $\text{KMnO}_4$ acts as a strong oxidant, and the $\text{Fe}^{2+}$ ions from Mohr's salt act as a reductant.

1. Preparation:

   • A standard solution of $\text{KMnO}_4$ (whose concentration is accurately known) is prepared and filled into a burette.
   • A known volume (e.g., $20.0 \text{ mL}$) of the unknown ferrous ammonium sulphate solution is accurately pipetted into a conical flask.
   • Dilute sulphuric acid is added to the conical flask to ensure the reaction occurs in a strongly acidic medium. This is crucial because $\text{MnO}_4^-$ is reduced to $\text{Mn}^{2+}$ (colourless) in acidic conditions, but to $\text{MnO}_2$ (brown precipitate) in neutral or alkaline conditions, which would obscure the end point.

2. Redox Reaction:

   The reaction involves the oxidation of $\text{Fe}^{2+}$ to $\text{Fe}^{3+}$ and the reduction of $\text{MnO}_4^-$ to $\text{Mn}^{2+}$ in an acidic medium.

   • Oxidation half-reaction: $\text{Fe}^{2+} \rightarrow \text{Fe}^{3+} + \text{e}^-$
   • Reduction half-reaction: $\text{MnO}_4^- + 8\text{H}^+ + 5\text{e}^- \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O}$

   To balance the electrons transferred, we multiply the oxidation half-reaction by 5:

   $5\text{Fe}^{2+} \rightarrow 5\text{Fe}^{3+} + 5\text{e}^-$

   Adding the two balanced half-reactions yields the overall balanced ionic equation:

   $\text{MnO}_4^- + 5\text{Fe}^{2+} + 8\text{H}^+ \rightarrow \text{Mn}^{2+} + 5\text{Fe}^{3+} + 4\text{H}_2\text{O}$

   Watch out

   Always ensure the medium is sufficiently acidic for permanganate titrations. If not, $\text{MnO}_4^-$ can be reduced to $\text{MnO}_2$, which is a brown precipitate, making the end point difficult to observe and leading to inaccurate results.

3. Titration Procedure:

   • The purple $\text{KMnO}_4$ solution is slowly added dropwise from the burette to the colourless ferrous ammonium sulphate solution in the conical flask. …