- The dissociation equation is \(H_2SO_4(aq) \rightleftharpoons 2H^+(aq) + SO_4^{2-}(aq)\).
- Adding \(HCl\) increases the concentration of the product \(H^+(aq)\).
- According to Le Chatelier's Principle, increasing the concentration of a product causes the system to shift towards the reactants to counteract the increase.
- The equilibrium shifts to the left, which corresponds to option 1. \(\boxed{1}\)
Explanation
- Analyzing the Equilibrium System
The chemical equilibrium for the dissociation of sulfuric acid is given by:
\[H_2SO_4(aq) \rightleftharpoons 2H^+(aq) + SO_4^{2-}(aq)\]
Sulfuric acid acts as a weak acid in its second dissociation step (though the first is strong), and we are looking at the general equilibrium state. In this equilibrium, \(H^+(aq)\) and \(SO_4^{2-}(aq)\) are the products, while \(H_2SO_4(aq)\) is the reactant.
- Effect of Adding Hydrochloric Acid
Hydrochloric acid (HCl) is a strong acid, meaning it dissociates completely in an aqueous solution according to the equation:
\[HCl(aq) \rightarrow H^+(aq) + Cl^-(aq)\]
When HCl is added to the system, it releases a significant amount of \(H^+(aq)\) ions into the solution. This addition of a product ion increases the total concentration of \(H^+(aq)\) in the equilibrium system.
- Applying Le Chatelier's Principle
According to Le Chatelier's Principle, if a system at chemical equilibrium is subjected to a change in concentration, temperature, or pressure, the system will adjust itself to counteract that change. Specifically, increasing the concentration of a product will cause the equilibrium to shift to the left, consuming the added species to return the system toward a new equilibrium state. Since the concentration of \(H^+(aq)\) has increased, the system will favor the reverse reaction to consume some of the excess \(H^+(aq)\) and form more \(H_2SO_4(aq)\). Therefore, the equilibrium shifts to the left.
- Conclusion
Based on the analysis, adding HCl increases the concentration of \(H^+(aq)\) on the product side. To counteract this, the system shifts to the left, consuming \(H^+\) and \(SO_4^{2-}\) to produce more \(H_2SO_4\). This matches option 1. Thus, the chemical equilibrium shifts to the left. The final answer is 1.
Examples
Le Chatelier's Principle is a fundamental concept used in industrial chemistry to maximize product yields. For instance, in the Haber process for producing ammonia (\(N_2 + 3H_2 \rightleftharpoons 2NH_3\)), manufacturers remove ammonia as it forms. By continuously removing a product, the equilibrium shifts to the right, forcing the system to produce more ammonia and increasing overall efficiency. This principle allows scientists to control chemical reactions effectively by manipulating concentrations, temperature, or pressure.