Problem 87 - Chemistry and Chemical Reactivity

Question

Describe the effect on the pH of the following actions or explain why there is not an effect: (a) Adding sodium acetate, \(\mathrm{NaCH}_{3} \mathrm{CO}_{2}\), to \(0.100 \mathrm{M} \mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}\) (b) Adding \(\mathrm{NaNO}_{3}\) to \(0.100 \mathrm{M} \mathrm{HNO}_{3}\)

Step-by-Step Solution

Answer

(a) Increases pH due to buffer formation; (b) No effect on pH.

Step 1: Determine the composition of sodium acetate solution

When sodium acetate, \( \mathrm{NaCH}_{3} \mathrm{CO}_{2} \), is added to the \( 0.100 \mathrm{M} \) acetic acid (\( \mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H} \)) solution, it dissociates into sodium ions \( \mathrm{Na}^{+} \) and acetate ions \( \mathrm{CH}_{3} \mathrm{CO}_{2}^{-} \). The acetate ion is a conjugate base of acetic acid.

Step 2: Analyze the buffer effect of acetate in acetic acid solution

The addition of acetate ions \( \mathrm{CH}_{3} \mathrm{CO}_{2}^{-} \) from sodium acetate creates a buffer solution by providing a common ion to the acetic acid. This buffer system helps to resist changes in pH when small amounts of acid or base are added, but also initially increases the pH, as the acetate ion can react with \( \mathrm{H}^{+} \).

Step 3: Assess the impact of NaNO3 addition to HNO3 solution

Adding \( \mathrm{NaNO}_{3} \) to the \( 0.100 \mathrm{M} \) \( \mathrm{HNO}_{3} \) solution results in the dissociation of \( \mathrm{NaNO}_{3} \) into \( \mathrm{Na}^{+} \) and \( \mathrm{NO}_{3}^{-} \) ions. However, since \( \mathrm{NO}_{3}^{-} \) is already present as part of \( \mathrm{HNO}_{3} \), no additional buffer or significant reaction occurs. The pH remains unaltered as \( \mathrm{HNO}_{3} \) is a strong acid and completely dissociates.

Key Concepts

Buffer Solutions

Buffer solutions are fascinating mixtures that help maintain a stable pH level even when small amounts of acids or bases are introduced. Essentially, a buffer solution combines a weak acid and its conjugate base, or a weak base and its conjugate acid. This duo has the unique ability to counteract pH changes. For example, when sodium acetate is added to an acetic acid solution, it forms a buffer. The acetate ions, which are the conjugate base of acetic acid, play a crucial role in this process. They "absorb" additional hydrogen ions (\(\mathrm{H}^+\)) from added acids. Similarly, when a base is added, the acetic acid can donate \(\mathrm{H}^+\), neutralizing the base.

Buffers are resilient, offering protection to their environment from drastic pH swings. This is particularly valuable in biological systems, where maintaining a consistent pH is critical for proper functioning. The stability of a buffer solution depends on its capacity, which is determined by the concentrations of the acid and its conjugate base. High concentrations provide a stronger buffer, capable of counteracting more substantial pH changes.

• Buffers involve a weak acid/base and its conjugate.
• Acid can neutralize added base; base can neutralize added acid.
• They are vital for stabilizing environments, notably in biology.

Common Ion Effect

The common ion effect is an intriguing principle of acid-base chemistry that explains the shift in equilibrium caused by the addition of an ion common to the solute. When we add sodium acetate to a solution of acetic acid, we introduce more acetate ions. These acetate ions are part of a common ion dynamic with the acetic acid already present.

According to Le Chatelier's principle, adding a common ion will shift the equilibrium to the left, reducing the ionization of the weak acid, and therefore, increasing the pH slightly. This is why, when sodium acetate, which dissociates into sodium and acetate ions, is introduced into an acetic acid solution, the pH rises. The acetate ions suppress the ionization of acetic acid by shifting the equilibrium towards the formation of the non-dissociated acid.

• The common ion effect reduces ionization of weak acids and bases.
• It shifts the equilibrium position to decrease ion products in solution.
• Le Chatelier's principle helps predict these shifts.

Acid Dissociation

Acid dissociation is the process by which an acid releases protons, forming its conjugate base. It is a fundamental concept that underlies many processes in chemistry. The extent of dissociation is quantified by the acid dissociation constant, \(K_a\), which measures the strength of an acid in solution. A high \(K_a\) value indicates a strong acid that dissociates completely, like nitric acid (HNO3), while a low \(K_a\) suggests a weaker acid, such as acetic acid (CH3CO2H), that dissociates only partially.

In the context of acid dissociation, when we add NaNO3 to an HNO3 solution, there is no notable effect on the pH. HNO3, being a strong acid, dissociates entirely in solution to release hydrogen and nitrate ions. The addition of NaNO3 does not change this equilibrium because it also releases nitrate ions, which do not affect the strong acid's complete dissociation. Thus, no shift in equilibrium occurs, and the pH remains steady.

• Acids dissociate to release protons, forming the conjugate base.
• \(K_a\) value assesses the strength of the acid.
• Strong acids fully dissociate, and no common ion effect alters this.