Question

Aqueous solutions of \(\mathrm{CoCl}_{2}\) are generally either light pink or blue. Low concentrations and low temperatures favor the pink form, whereas high concentrations and high temperatures favor the blue form. Adding hydrochloric acid to a pink solution of \(\mathrm{CoCl}_{2}\) causes the solution to turn blue; the pink color is restored by the addition of \(\mathrm{HgCl}_{2}\). Account for these observations.

Short answer

Cl^- concentration and equilibrium shifts explain the color change.

Step-by-step solution

Understanding the Chemical Reaction

When CoCl_2 is dissolved in water, it can exist in equilibrium between different forms, generally pink ( Co(H_2O)_6^{2+}) and blue ( CoCl_4^{2-}). The reaction can be represented as: Co(H_2O)_6^{2+} (pink) + 4Cl^- <=> CoCl_4^{2-} (blue) + 6H_2O.

Effect of Concentration and Temperature

Higher concentrations of Cl^- ions and higher temperatures shift the equilibrium towards the blue CoCl_4^{2-} form, while lower concentrations and temperatures favor the pink Co(H_2O)_6^{2+} form.

Effect of Adding Hydrochloric Acid

Adding hydrochloric acid, HCl, increases the concentration of Cl^- ions in the solution. According to Le Chatelier's principle, this will shift the equilibrium towards the blue CoCl_4^{2-} form.

Effect of Adding HgCl_2

Adding HgCl_2 introduces Hg^{2+} ions, which form a stable complex with Cl^- ions ( HgCl_4^{2-}). This reduces the concentration of free Cl^- ions in the solution, shifting the equilibrium back towards the pink Co(H_2O)_6^{2+} form.

Key concepts

Chemical Equilibrium

A chemical equilibrium occurs when a chemical reaction reaches a state where the concentrations of reactants and products remain constant over time. This does not mean the reaction stops but that the forward and reverse reactions occur at the same rate. For the \(\mathrm{CoCl}_2\) aqueous solution, the equilibrium lies between two different forms of cobalt complexes: \(\text{pink } \mathrm{Co}(\mathrm{H}_2\mathrm{O})_6^{2+} \) and \(\text{blue } \mathrm{CoCl}_4^{2-}\).

Understanding the position of equilibrium is essential. This position can shift based on various factors such as concentration and temperature. However, the crucial aspect of equilibrium is that both reactants and products are present, allowing for dynamic changes when the system is disturbed.

Le Chatelier's Principle

Le Chatelier's principle helps predict how a change in conditions affects a chemical equilibrium. The principle states that if a system at equilibrium experiences a change in concentration, temperature, or pressure, the equilibrium will shift to counteract this change.

When hydrochloric acid (\(\mathrm{HCl}\)) is added to our \(\mathrm{CoCl}_2\) solution, the concentration of \(\mathrm{Cl}^-\) ions increases. According to Le Chatelier's principle, the equilibrium shifts to the right to form more \(\text{blue } \mathrm{CoCl}_4^{2-}\), absorbing the excess chloride ions.
This shift results in the observed color change from pink to blue. It visually represents the system's response to restore equilibrium by balancing the stress introduced.

Effect of Temperature and Concentration

Concentration and temperature are pivotal factors affecting chemical equilibrium. In the case of \(\mathrm{CoCl}_2\) solutions:

• Lower concentrations of \(\mathrm{Cl}^-\) ions and reduced temperatures favor the pink \(\mathrm{Co}(\mathrm{H}_2\mathrm{O})_6^{2+}\) complex.
• Higher concentrations of \(\mathrm{Cl}^-\) ions due to \(\mathrm{HCl}\) addition, or increased temperatures, favor the blue \(\mathrm{CoCl}_4^{2-}\) complex.

When you alter the concentration by adding a strong electrolyte like \(\mathrm{HCl}\), the equilibrium tends to shift to balance the new concentration of ions. Similarly, increasing the temperature can drive the equilibrium toward the endothermic side.
It's crucial for students to comprehend how both of these changes act to shift the equilibrium and influence the overall color of the solution.

Complex Ion Formation

In aqueous chemistry, the formation of complex ions can greatly influence the properties and reactions of solutions. A complex ion consists of a central metal ion bonded to molecules or ions. These formations can alter equilibrium positions.

For example, when \(\mathrm{HgCl}_2\) is added to the \(\mathrm{CoCl}_2\) solution, \(\mathrm{Hg}^{2+}\) ions form a stable complex with chloride ions, specifically \(\mathrm{HgCl}_4^{2-}\). This reduces the free \(\mathrm{Cl}^-\) ion concentration, causing the equilibrium to shift back towards the pink \(\mathrm{Co}(\mathrm{H}_2\mathrm{O})_6^{2+}\) complex.

• Complex forming ions can "tie-up" ligands (like \(\mathrm{Cl}^-\)), reducing their availability.
• The shift in equilibrium shows a clear interplay of complexation competing with the existing equilibrium, visible through color changes in the solution.

Recognizing and understanding complex ion formation is key to mastering the chemistry of solutions.