The equilibrium reaction given is: $2{\mathrm{Cu}}^I\rightleftharpoons {\mathrm{Cu}}^0+{\mathrm{Cu}}^{II}$. This represents the disproportionation of Cu⁺ (copper(I)) into Cu⁰ (copper metal) and Cu²⁺ (copper(II)).

For a disproportionation reaction to be spontaneous, the product species must be stabilized. The direction of the equilibrium can be shifted by ligands that form stable complexes with specific oxidation states of copper, thereby lowering their energy and favoring the side of the reaction where the complex is formed.

In aqueous medium, Cu⁺ is unstable and disproportionates. However, the presence of certain ligands can stabilize Cu⁺ by forming complexes, which shifts the equilibrium to the left (towards Cu⁺). The ligand that forms the most stable complex with Cu⁺ will best prevent its disproportionation.

Among the given ligands:

• CN⁻ forms a very stable complex with Cu⁺, [Cu(CN)₂]⁻ or [Cu(CN)₄]³⁻, which has a high formation constant. This strongly stabilizes Cu⁺ and shifts the equilibrium left.
• Cl⁻ forms complexes with both Cu⁺ (e.g., [CuCl₂]⁻) and Cu²⁺ (e.g., [CuCl₄]²⁻), but the complex with Cu⁺ is not exceptionally stable, so it does not shift the equilibrium as effectively.
• NO₃⁻ is a very weak ligand and does not form significant complexes with copper ions; it has minimal effect on the equilibrium.
• SCN⁻ (thiocyanate) can coordinate through S or N. With Cu⁺, it may form complexes, but they are less stable than cyanide complexes. It is not as effective at stabilizing Cu⁺.

Therefore, CN⁻ is the ligand that most effectively shifts the equilibrium towards the left due to the high stability of the cyanide complex with Cu⁺.

Related Topics & Formulae

Complex Formation and Stability Constants: The stability constant (K) quantifies the stability of a complex. A higher K value means a more stable complex. For Cu⁺ with CN⁻, the stability constant is very high, which explains the shift in equilibrium.

Disproportionation: A redox reaction where a single species is both oxidized and reduced. The tendency can be predicted using standard reduction potentials. For Cu⁺, E° for Cu⁺ + e⁻ → Cu⁰ is +0.52 V and E° for Cu²⁺ + e⁻ → Cu⁺ is +0.15 V. Since E°_reduction for the reduction half is greater than E°_oxidation for the oxidation half, disproportionation is favorable (E°_cell = +0.37 V > 0). Ligands change the effective reduction potentials by complex formation.