Introduction

Colorimetry is a powerful analytical technique that uses the absorption of light by colored solutions to determine chemical concentrations. In HSC Chemistry, this method is particularly valuable for studying equilibrium systems, especially the iron(III) thiocyanate equilibrium.

The Iron(III) Thiocyanate Equilibrium System

The formation of the iron(III) thiocyanate complex involves a reversible reaction between iron(III) ions and thiocyanate ions:

${Fe^{3+}(aq) + SCN^-(aq) <=> FeSCN^{2+}(aq)}$

This reaction produces a distinctive deep red complex, making it ideal for colorimetric analysis.

Understanding Beer-Lambert Law

The fundamental principle behind colorimetry is the Beer-Lambert Law:

$A = \varepsilon c l$

Where:

Experimental Procedure

1. Preparation of Standard Solutions

2. Absorbance Measurements

3. Creating the Calibration Curve

Plot absorbance versus concentration of FeSCN^{2+}. This creates a linear relationship following Beer-Lambert Law.

4. Equilibrium Analysis

After measuring the equilibrium mixture's absorbance, calculate the equilibrium concentrations:

$[Fe^{3+}]_{eq} = [Fe^{3+}]_{initial} - [FeSCN^{2+}]_{eq}$ $[SCN^-]_{eq} = [SCN^-]_{initial} - [FeSCN^{2+}]_{eq}$

5. Calculating K_{eq}

The equilibrium constant is determined using:

$K_{eq} = \frac{[FeSCN^{2+}]_{eq}}{[Fe^{3+}]_{eq} \times [SCN^-]_{eq}}$