- The reaction is a heterogeneous equilibrium governed by the solubility product constant, \(K_{sp} = [Ag^{+}][Cl^{-}]\).
- Adding \(AgNO_3\) significantly increases the concentration of \(Ag^{+}\) ions because it is a strong electrolyte.
- According to Le Chatelier's principle, the system compensates for this increase by shifting to consume the excess silver ions.
- This shift occurs by moving the equilibrium to the left, resulting in the formation of more solid \(AgCl\).
- The chemical equilibrium of the system shifts to the left. \(\boxed{\text{The chemical equilibrium of the system shifts to the left.}}\)
Explanation
- Analysis of the Chemical Equilibrium
The reaction provided represents the dissolution of silver chloride (AgCl) in water, a heterogeneous equilibrium described by:
\[AgCl(s) \rightleftharpoons Ag^{+}(aq) + Cl^{-}(aq)\]
In this system, the equilibrium is characterized by the solubility product constant, \(K_{sp}\), defined as:
\[K_{sp} = [Ag^{+}][Cl^{-}]\]
This expression indicates the product of the concentrations of silver ions (\(Ag^{+}\)) and chloride ions (\(Cl^{-}\)) at equilibrium. The state of this equilibrium depends on the concentration of the dissolved ions.
- Disturbance of the Equilibrium by Silver Nitrate Addition
When silver nitrate (\(AgNO_3\)) is added to the system, it acts as a strong electrolyte and fully dissociates in water according to the following equation:
\[AgNO_3(aq) \rightarrow Ag^{+}(aq) + NO_3^{-}(aq)\]
This process increases the concentration of silver ions, \([Ag^{+}]\), in the solution. According to Le Chatelier's principle, if a system at equilibrium is disturbed by a change in concentration, the system will adjust itself to counteract that disturbance. Since \([Ag^{+}]\) has increased, the equilibrium will shift to consume some of the excess silver ions.
- Determining the Shift of the Equilibrium
To consume the additional \(Ag^{+}\) ions, the chemical system shifts in the direction that removes these ions from the solution. Looking at the original equilibrium equation:
\[AgCl(s) \rightleftharpoons Ag^{+}(aq) + Cl^{-}(aq)\]
The reverse reaction combines \(Ag^{+}\) and \(Cl^{-}\) ions to form the solid precipitate \(AgCl(s)\). Therefore, the addition of \(AgNO_3\) forces the equilibrium to shift to the left, which results in the precipitation of more solid \(AgCl\). Consequently, the solubility of silver chloride decreases in the presence of an excess of silver ions, a phenomenon known as the common-ion effect.
Examples
The common-ion effect is frequently utilized in industrial and laboratory settings for precipitation processes. For example, when purifying silver compounds, chemists can intentionally add a substance containing common ions to force the precipitation of a target salt. This is a practical application of manipulating equilibrium to control chemical purity or yield in manufacturing processes, demonstrating how understanding chemical principles directly impacts efficiency and control in real-world chemical production.