Question

Which one of the following is not a buffer solution? (a) \(0.8 \mathrm{M} \mathrm{H}_{2} \mathrm{~S}+0.8 \mathrm{M} \mathrm{KHS}\) (b) \(2 \mathrm{M} \mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{2}+2 \mathrm{M} \mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{3}^{+} \mathrm{Br}^{-}\) (c) \(3 \mathrm{M} \mathrm{H}_{2} \mathrm{CO}_{3}+3 \mathrm{M} \mathrm{KHCO}_{3}\) (d) \(0.05 \mathrm{M} \mathrm{KCIO}_{4}+0.05 \mathrm{M} \mathrm{HClO}_{4}\)

Short answer

Option (d) is not a buffer solution.

Step-by-step solution

Understand the Buffer System

A buffer solution contains a weak acid with its conjugate base or a weak base with its conjugate acid. It functions to maintain a relatively constant pH when small amounts of acid or base are added.

Analyze Option (a)

The given components are a weak acid (6 H_2S) and its conjugate base (KHS). This forms a buffer system.

Analyze Option (b)

Here, we have a weak base (6 C_6H_5NH_2) and its conjugate acid (C_6H_5NH_3^+). This combination constitutes a buffer solution.

Analyze Option (c)

The solution includes a weak acid (6 H_2CO_3) and its conjugate base (KHCO_3). This is a typical buffer solution.

Analyze Option (d)

This combination contains a salt (KClO_4) and a strong acid (6 HClO_4). Since HClO_4 is a strong acid, it does not pair with a suitable weak base to form a buffer.

Determine the Non-buffer Solution

Option (d) does not satisfy the requirement for a buffer system because it includes a strong acid without a conjugate base pair. A buffer requires a weak acid-base combination.

Key concepts

Weak Acid and Conjugate Base

Buffer solutions rely on the presence of a weak acid and its conjugate base to function effectively. Let's delve into this concept: a weak acid only partially dissociates in water, remaining mostly intact while releasing a small number of hydrogen ions (H⁺). A common example of this is acetic acid (CH₃COOH). The conjugate base, on the other hand, is formed when the weak acid loses a proton (H⁺). In the case of acetic acid, the conjugate base is the acetate ion (CH₃COO⁻). Together, they work to neutralize any added strong base or acid.

• The weak acid is in equilibrium with its dissociation products.
• The conjugate base reacts with added acids, mitigating pH changes.

This partnership stabilizes the pH of a solution by reacting with any acids or bases added, thus keeping the solution's pH within a narrow range.

Weak Base and Conjugate Acid

Similarly, buffer solutions can be composed of a weak base and its conjugate acid. Let's consider this setup: a weak base typically accepts protons (H⁺) and is characterized by its limited ionization in water. Ammonia (NH₃) is a classic example of a weak base. Upon gaining a proton, a weak base forms its conjugate acid. For ammonia, the conjugate acid is the ammonium ion (NH₄⁺). This conjugate acid acts as a proton donor when necessary, complementing the behavior of the weak base.

• The weak base is in equilibrium with its reacting products in solution.
• The conjugate acid can neutralize any bases added to the solution, preventing pH shifts.

This balanced interaction between the base and its conjugate acid prevents drastic pH changes, ensuring the buffer's effectiveness.

pH Stability

The primary function of a buffer solution is to maintain pH stability. This stability is crucial in biological systems, where enzymes and other biochemical processes depend on consistent pH levels to function optimally. Here's how buffer solutions achieve pH stability: When an acid is added to a buffer solution, the conjugate base present in the solution neutralizes the extra hydrogen ions (H⁺), preventing a significant change in pH. Similarly, when a base is added, the weak acid converts the extra hydroxide ions (OH⁻) into water, minimizing any increase in pH.

• Buffers can resist changes in pH upon small additions of acid or base.
• They are essential in biological and chemical processes where pH control is necessary.

This resistance to pH change is what allows buffers to stabilize environments, making them integral in many scientific applications.

Neutralization Reactions

Neutralization reactions are a fundamental part of how buffer solutions control pH. These reactions involve the direct interaction of acids with bases to form either more neutral or less reactive species. In a buffer, when an added acid attempts to lower the pH, the conjugate base present in the solution will engage in a neutralization reaction. This reaction will produce water and other compounds, often leaving the pH unchanged. Conversely, when a base is introduced to the buffer, the weak acid in the solution takes on the role of neutralizing the base, engaging in another neutralization reaction that again stabilizes the pH.

• Neutralization reactions convert strong acids/bases to weaker ones or water.
• They are key to the buffer's ability to maintain a consistent pH environment.

Understanding these reactions is crucial as they highlight how buffers operate when challenged by pH-altering substances.