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The common-ion effect occurs when a compound is added to a solution in equilibrium and that compound contains an ion already present in the solution. In our scenario, the solution contains methylamine \((\text{CH}_3\text{NH}_2)\), which partially ionizes to produce \(\text{CH}_3\text{NH}_3^+\) ions and hydroxide \((\text{OH}^-)\) ions. Adding \(\text{CH}_3\text{NH}_3\text{Cl} \), which dissociates into \(\text{CH}_3\text{NH}_3^+\) and \(\text{Cl}^-\), increases the concentration of the common ion, \(\text{CH}_3\text{NH}_3^+\).

This additional \(\text{CH}_3\text{NH}_3^+\) exerts the common-ion effect by shifting the equilibrium of the methylamine ionization reaction. This causes changes in the concentrations of the products and reactants involved.

Methylamine \((\text{CH}_3\text{NH}_2)\) is classified as a weak base. Unlike strong bases, weak bases do not fully ionize in water. Instead, they establish an equilibrium between the non-ionized and ionized forms. In the case of methylamine, the equilibrium can be represented by the equation:

\( \text{CH}_3\text{NH}_2 + \text{H}_2\text{O} \rightleftharpoons \text{CH}_3\text{NH}_3^+ + \text{OH}^- \).

Since it is a weak base, only a fraction of the methylamine molecules ionize in water, leading to a solution that is basic but not as strongly basic as it would be with a strong base. This partial ionization is critical for the behavior observed when common ions are introduced.

The concept of equilibrium shift is vital here. When a change is applied to a system at equilibrium, the system responds to counter the effect of that change, as per Le Chatelier's principle. By adding \(\text{CH}_3\text{NH}_3\text{Cl}\) to the methylamine solution, the concentration of \(\text{CH}_3\text{NH}_3^+\) increases.

According to Le Chatelier's principle, the equilibrium will shift to oppose the change, in this case by decreasing \(\text{CH}_3\text{NH}_3^+\) concentration.

• The reaction shifts to the left, indicating a decrease in the forward reaction where \(\text{CH}_3\text{NH}_2\) ionizes into \(\text{CH}_3\text{NH}_3^+\) and \(\text{OH}^-\).
• This shift results in fewer \(\text{OH}^-\) ions, meaning the system is becoming less basic.

The pH change in this experiment is due to the equilibrium shift caused by the common-ion effect. Initially, the solution of \(\text{CH}_3\text{NH}_2\) has a high pH of 11.8, indicating basicity. When \(\text{CH}_3\text{NH}_3\text{Cl}\) is added, the concentration of hydroxide ions \((\text{OH}^-)\) diminishes as the equilibrium shifts to the left.

With fewer \(\text{OH}^-\) ions in the solution, the basicity decreases, resulting in a lower pH. This change demonstrates how pH reflects the balance of hydrogen \((\text{H}^+)\) and hydroxide \((\text{OH}^-)\) ions in a solution and how it can be affected by adding substances that alter this balance.

• The decrease in \(\text{OH}^-\) reflects a smaller base strength.
• This leads to a decrease in pH, showcased by the system moving towards neutrality.