Question

Explain, with balanced ionic equations, why (a) \(\mathrm{CuI}_{2}\) dissolves in ammonia solution, (b) AgBr dissolves in NaCN solution, (c) \(\mathrm{Hg}_{2} \mathrm{Cl}_{2}\) dissolves in \(\mathrm{KCl}\) solution.

Short answer

CuI2, AgBr and Hg2Cl2 dissolve in NH3, NaCN and KCl solutions due to the formation of complex ions. The balanced ionic equations are: For CuI2 in NH3 : CuI2 + 4NH3 -> [Cu(NH3)4]2+ + 2I-. For AgBr in NaCN : AgBr + 2CN- -> [Ag(CN)2]- + Br-. For Hg2Cl2 in KCl : Hg2Cl2 + 2Cl- -> [HgCl2]2+ + 2Cl-.

Step-by-step solution

Dissolution of CuI2 in Ammonia solution

When \(\mathrm{CuI}_{2}\) dissolves in ammonia solution, the equation is expressed as follows: \(\mathrm{CuI}_{2} + 4NH_{3}\) \( \rightarrow \) \([Cu(NH_{3})_{4}]^{2+} + 2I^{-}\) Here, the \(\mathrm{CuI}_{2}\) ionizes in the ammonia solution to form a complex ion and iodine ion.

Dissolution of AgBr in NaCN solution

When AgBr dissolves in NaCN solution, the ionic equation is written as follows: \(\mathrm{AgBr} +2CN^{-}\) \( \rightarrow \) \([Ag(CN)_{2}]^{-} + Br^{-}\) Here, the \(\mathrm{AgBr}\) ionizes to form a complex ion and bromine ion.

Dissolution of Hg2Cl2 in KCl solution

When \(\mathrm{Hg}_{2} \mathrm{Cl}_{2}\) dissolves in \(\mathrm{KCl}\) solution, the equation is expressed as follows: \(\mathrm{Hg}_{2} \mathrm{Cl}_{2} + 2Cl^{-}\) \( \rightarrow \) \([HgCl_{2}]^{2+} + 2Cl^{-}\) Here, the \(\mathrm{Hg}_{2} \mathrm{Cl}_{2}\) ionizes in the \(\mathrm{KCl}\) solution to form a complex ion and chloride ion.

Key concepts

Complex Ion Formation

Complex ion formation is a fascinating chemical process where molecules or ions bond to a central metal ion, creating a larger, more complex molecule. These complexes often form when a metal ion combines with ligands. Ligands are neutral molecules or ions with one or more pairs of electrons that can be donated to the metal ion. This process is crucial in understanding how some substances dissolve in solution.
For example, when \( \mathrm{CuI}_2 \) interacts with ammonia, it forms a complex ion as \( \mathrm{CuI}_2 + 4NH_3 \rightarrow [Cu(NH_3)_4]^{2+} + 2I^- \).Here, the ammonia molecules, acting as ligands, attach to the copper ion. The resulting complex ion is more soluble than the initial compound, leading to the dissolution of \( \mathrm{CuI}_2 \). This principle of complex ion formation helps explain why certain sparingly soluble compounds become more soluble in appropriate solvents.

Dissolution Reaction

Dissolution reactions describe the process where a solid substance dissolves in a solvent, forming a solution. This involves the breakdown of ionic compounds into their constituent ions, which then disperse throughout the solvent. In our context, dissolution is facilitated by creating complex ions.
Take the example of \( \mathrm{AgBr} \) dissolving in\( \mathrm{NaCN} \) solution. The reaction goes as follows: \( \mathrm{AgBr} + 2CN^- \rightarrow [Ag(CN)_2]^- + Br^- \).In this reaction, silver bromide, which is relatively insoluble in water, dissolves because the formation of the complex ion \([Ag(CN)_2]^-\) makes the entire compound more soluble. This happens because the complex ion is more stable than the separate ions, driving the reaction towards dissolution.

Coordination Chemistry

Coordination chemistry is the study of complex ions, focusing on how metal ions interact with ligands. This field explains why metals can bond with various molecules to form stable complexes, influencing solubility and reactivity.
For instance, when \( \mathrm{Hg}_2\mathrm{Cl}_2 \) dissolves in \( \mathrm{KCl} \),the following process occurs: \( \mathrm{Hg}_2\mathrm{Cl}_2 + 2Cl^- \rightarrow [HgCl_2]^{2+} + 2Cl^- \).Here, excess chloride ions from \( \mathrm{KCl} \) coordinate with mercury to form a complex. This process enhances the solubility of \( \mathrm{Hg}_2\mathrm{Cl}_2 \), which is otherwise insoluble under normal conditions. Coordination chemistry provides insight into how these interactions lead to the formation of more stable or more soluble compounds, thereby facilitating dissolution.