Question

Use Le Châtelier's principle to explain how the common ion effect affects the \(\mathrm{pH}\) of a weak acid solution.

Short answer

The common ion effect increases the pH of a weak acid solution by reducing hydronium ion concentration.

Step-by-step solution

Understanding the Common Ion Effect

The common ion effect occurs when a compound containing an ion already present in a solution is added to the solution. This additional ion can shift the equilibrium position according to Le Châtelier's principle, affecting the solubility or dissociation of weak acids or bases.

Establishing the Initial Equilibrium

Consider a weak acid HA which partially dissociates in water, following the equilibrium \[ ext{HA} ightleftharpoons ext{H}^+ + ext{A}^- \]Here, HA dissociates to form H+ (the hydrogen ion) and A- (the conjugate base).

Analyzing the Addition of Common Ion

When a salt containing a common ion A-, such as sodium acetate \(( ext{NaA})\), is added to the solution, the concentration of A- increases. This perturbs the equilibrium set by the weak acid.

Applying Le Châtelier's Principle

According to Le Châtelier's principle, adding more A- shifts the equilibrium to the left. This means the reaction will favor the formation of more HA and consume some H+ ions to reduce the stress of added A-.

Effect on pH

Since the equilibrium shift results in fewer free hydrogen ions \(( ext{H}^+)\), the solution becomes less acidic, leading to an increase in the pH of the solution. This demonstrates the common ion effect.

Key concepts

Common Ion Effect

When a solution has a compound added that shares an ion with substances already in the solution, it triggers what is known as the "common ion effect." Imagine you have a solution with a weak acid and you add a salt containing the same anion as the acid's conjugate base. This additional anion increases the concentration of common ions in the solution.
According to Le Châtelier's principle, the equilibrium of the weak acid will shift to counteract this change. This shift typically means that the dissociation of the weak acid is suppressed, which affects both the concentration of hydrogen ions and the acidity of the solution.

Weak Acid

A weak acid is a type of acid that does not fully dissociate, or break apart, into ions in a solution. Unlike strong acids, which completely ionize, weak acids establish a balance between their undissociated form and ions.
For instance, when acetic acid (a common example of a weak acid) is dissolved in water, only a small fraction of acetic acid molecules will dissociate into hydrogen ion (H^+) and acetate ion (A^-). This balance can be represented as: \[ \text{HA} \rightleftharpoons \text{H}^+ + \text{A}^- \]Because only a portion of the molecules dissociate, the equilibrium between the molecules that remain undissociated and the ions is crucial for understanding how weak acids behave in different conditions.

Equilibrium

In chemistry, equilibrium refers to the state where the rate of the forward reaction equals the rate of the reverse reaction. For a weak acid in solution, equilibrium is what maintains the balance between the acid molecules and their ions.
Using Le Châtelier's principle, we can predict how equilibrium will respond to changes. For weak acids, adding more of the conjugate base from an external source, like a salt containing the base, will cause the equilibrium to shift. This shift reduces the number of free hydrogen ions and in turn decreases the solution's acidity, thereby increasing the pH.

pH

pH measures how acidic or basic a solution is. It's calculated based on the concentration of hydrogen ions ( H^+ ) in the solution. The more hydrogen ions present, the more acidic the solution and the lower the pH value.
With weak acids, the presence of a common ion can significantly affect the pH. When the concentration of the hydrogen ions decreases, due to the common ion effect shifting the equilibrium, the pH increases.

• This rise in pH reflects the solution becoming less acidic.
• Understanding this relationship helps in predicting the behavior of weak acids when exposed to additional ions.

As such, the interplay between weak acids, common ions, and Le Châtelier's principle is crucial for understanding changes in solution acidity.