In equilibrium: \(\mathrm{SCN}(\mathrm{aq})+\mathrm{Fe}^{+3}(\mathrm{aq}) \rightleftharpoons\left[\mathrm{Fe}(\mathrm{SCN})^{2+}\right](\mathrm{aq})\) \(\begin{array}{lll}\text { colourless } & \text { yellow } & \text { deep red }\end{array}\) If thiocyanate ions are added in equilibrium mixture (a) The solution becomes colourless. (b) The yellow colour of solution deepens. (c) The red colour of the solution deepens. (d) Concentration of \([\mathrm{Fe}(\mathrm{SCN})]^{2+}\) ion will decrease.

Le Chatelier's principle is a fundamental concept in chemical equilibrium that describes how a system at equilibrium reacts to disturbances. This principle states that if an external change is applied to a system in equilibrium, the system will adjust itself to counteract that change and re-establish equilibrium. Such changes can include variations in concentration of reactants or products, changes in pressure or volume for gaseous systems, or alterations in temperature.

For example, when the concentration of a reactant is increased, the equilibrium will shift towards the side of the reaction that will use up the added reactant, thereby producing more products. If the concentration of a product is increased, the system will shift the other way. The principle helps us predict the outcomes of changing conditions on the concentrations of reactants and products within a chemical reaction.

The equilibrium constant, denoted as K, quantifies the balance between the concentrations of reactants and products in a chemical reaction at equilibrium. It is a fixed value at a given temperature and is defined by the concentrations (or pressures for gases) of the reactants and products raised to the power of their coefficients from the balanced equation. For the general reaction \(aA + bB \rightleftharpoons cC + dD\), the equilibrium constant \(K_{c}\) is given by \(K_{c} = \frac{[C]^{c}[D]^{d}}{[A]^{a}[B]^{b}}\).

In the context of our exercise, if the concentration of SCN (thiocyanate ions) is increased, according to Le Chatelier's principle, the equilibrium position will shift to the right to form more of the complex iron(III) thiocyanate, \( [Fe(SCN)]^{2+} \), which means that the actual concentration ratio of the products to reactants will change. However, the equilibrium constant itself remains unchanged, as it is only dependent on the temperature.

Complex ion formation involves the combination of a metal ion with other molecules or ions, often called ligands, resulting in a species known as a complex ion. These complex ions usually have distinct colors based on the metal and ligands involved, offering a visual cue of their presence in a solution. In our exercise, the metal ion \(Fe^{+3}\) reacts with the ligand \(SCN^-\) to form the deep red complex ion \( [Fe(SCN)]^{2+} \).

The formation of a complex ion can significantly alter the equilibrium position of a reaction. Adding more ligands (like SCN-) to the equilibrium mixture leads to an increased formation of the complex ion, which is a vivid example of Le Chatelier's principle at work. The deepening of the red color indicates more complex ion is formed and also showcases the dynamic nature of equilibrium. Complex ion formation is a key concept in coordination chemistry and is exploited in various analytical and industrial processes.