- The added \(MgCl_2\) dissociates into \(Mg^{2+}\) and \(Cl^{-}\) ions, which do not participate in the equilibrium reaction of acetic acid.
- Since magnesium acetate is soluble and hydrochloric acid is a strong acid that remains dissociated, the ions do not react with \(CH_3COO^{-}\) or \(H^{+}\).
- Because no reactants or products are consumed or produced by the added ions, the equilibrium position remains unchanged according to Le Chatelier's principle.
- The final answer is \(\boxed{\text{The chemical equilibrium of the system will be unaffected.}}\)
Explanation
- Analysis of the System and Added Compound
The chemical equilibrium in question is:
\[CH_3COOH(aq) \rightleftharpoons CH_3COO^{-}(aq) + H^{+}(aq)\]
We are asked to determine the effect on this equilibrium when magnesium chloride (\(MgCl_2\)) is added to the solution. First, we note that \(MgCl_2\) is a salt that dissociates completely in water according to the following equation:
\[MgCl_2(s) \rightarrow Mg^{2+}(aq) + 2Cl^{-}(aq)\]
The resulting ions, \(Mg^{2+}\) and \(Cl^{-}\), are introduced into the equilibrium mixture. To determine if this affects the position of the equilibrium, we must evaluate whether these ions interact with the components of the acetic acid equilibrium (\(CH_3COOH\), \(CH_3COO^{-}\), or \(H^{+}\)).
- Evaluating Potential Interactions and Equilibrium Shift
Next, we examine the possible reactions between the added ions and the existing species:
- The acetate ion (\(CH_3COO^{-}\)) would form magnesium acetate if it precipitated with \(Mg^{2+}\). However, magnesium acetate is highly soluble in water. Thus, no precipitate forms, and the concentration of \(CH_3COO^{-}\) remains essentially unchanged.
- The hydrogen ion (\(H^{+}\)) would form hydrogen chloride if it reacted with \(Cl^{-}\). However, HCl is a strong acid that exists in solution entirely in its dissociated form (\(H^{+}\) and \(Cl^{-}\)). Therefore, adding \(Cl^{-}\) does not remove \(H^{+}\) from the system.
Since neither the magnesium ions nor the chloride ions effectively consume or produce any of the species participating in the acetic acid dissociation equilibrium, they act as spectator ions. Their presence does not change the concentrations of \(CH_3COOH\), \(CH_3COO^{-}\), or \(H^{+}\) in a way that would shift the equilibrium position based on Le Chatelier's principle.
- Conclusion
Because the added \(Mg^{2+}\) and \(Cl^{-}\) ions do not react with the components of the acetic acid equilibrium, the state of the equilibrium remains constant. There is no stress applied to the system that would favor either the forward or the reverse reaction. Therefore, we conclude that the chemical equilibrium of the system will be unaffected. The final answer is that the chemical equilibrium of the system will be unaffected.
Examples
In chemistry laboratories, identifying whether a salt is a spectator ion or a reactant is crucial for understanding reaction mechanisms. For instance, when adding salts like KCl or NaCl to various equilibrium systems, chemists check if the ions are inert or involved in precipitation or acid-base reactions. This prevents unnecessary modifications to the desired chemical output and ensures experiment accuracy. Understanding this helps in predicting whether adding an ionic compound will alter the pH or reaction yield of your chemical setup.