Question

Which of the following gives a buffer solution when equal volumes of the two solutions are mixed? (a) \(0.10 \mathrm{M} \mathrm{HF}\) and \(0.10 \mathrm{M} \mathrm{NaF}\) (b) \(0.10 \mathrm{M} \mathrm{HF}\) and \(0.10 \mathrm{M} \mathrm{NaOH}\) (c) \(0.20 \mathrm{M} \mathrm{HF}\) and \(0.10 \mathrm{M} \mathrm{NaOH}\) (d) \(0.10 \mathrm{M} \mathrm{HCl}\) and \(0.20 \mathrm{M} \mathrm{NaF}\)

Short answer

Options (a) and (c) form buffer solutions.

Step-by-step solution

Identify the Requirements for a Buffer

A buffer solution consists of a weak acid and its conjugate base or a weak base and its conjugate acid. It can also be made from a weak acid and a strong base in the correct proportions or vice versa.

Analyze Option (a)

Option (a) consists of 0.10 M HF and 0.10 M NaF. HF is a weak acid and NaF provides the conjugate base F⁻. This forms a buffer solution since it contains a weak acid and its conjugate base.

Analyze Option (b)

Option (b) consists of 0.10 M HF and 0.10 M NaOH. Upon mixing, HF will react with NaOH, a strong base, to form F⁻, but all HF will be neutralized. Therefore, this does not form a buffer solution as the weak acid is completely consumed.

Analyze Option (c)

Option (c) consists of 0.20 M HF and 0.10 M NaOH. Mixing them in equal volumes will result in HF partially reacting with NaOH, forming some F⁻ and leaving some HF unreacted. This creates a mixture of weak acid (HF) and its conjugate base (F⁻), forming a buffer solution.

Analyze Option (d)

Option (d) consists of 0.10 M HCl and 0.20 M NaF. HCl is a strong acid which will fully dissociate, leaving no weak acid. Therefore, this does not form a buffer solution as there is no weak acid/base involved.

Determine the Correct Answer

Only options (a) and (c) satisfy the conditions required for forming a buffer solution: a weak acid and its conjugate base in suitable proportions.

Key concepts

Weak Acid

A weak acid is an acid that does not completely dissociate into its ions in water. This means only a small fraction of the acid molecules will split into hydrogen ions ( H^+ ) and their corresponding negatively charged ions. The most critical property of a weak acid is its equilibrium state in solution, which allows for partial dissociation.

Weak acids play a significant role in forming buffer solutions. These solutions resist changes in pH when small amounts of acids or bases are added. In the context of a buffer solution, a weak acid provides the crucial component. It ensures that upon slight addition of an acid or base, the pH of the solution remains relatively constant.

For example, hydrofluoric acid ( HF ) used in the exercise above is a weak acid. In a buffer, it ensures stable pH levels because it doesn’t fully dissociate in water. This partial dissociation is precisely what provides buffers their capability to protect against pH changes.

Conjugate Base

The conjugate base of an acid is what remains after the acid has donated a hydrogen ion ( H^+ ). When a weak acid, like HF, loses an H^+ , it becomes its conjugate base, F^- .

Conjugate bases are crucial in creating buffer solutions. This is because they can accept hydrogen ions, helping to balance any changes in hydrogen ion concentration that occur due to the addition of acids or bases. Thus, they maintain the pH balance of the solution.

In buffer solutions, typically, you have a weak acid and its conjugate base. This pair works together to absorb any hydrogen or hydroxide ions introduced, thereby preventing drastic shifts in pH. For example, in the problem, NaF provides the F^- , the conjugate base of HF , enabling the buffer action.

Chemical Equilibrium

Chemical equilibrium refers to the state where the rates of the forward and reverse reactions are equal. In other words, the concentration of reactants and products remains constant over time.

In the context of buffers and weak acids, chemical equilibrium is vital. It ensures that there is a constant exchange between the acid and its conjugate base. When a weak acid is in equilibrium, it partially dissociates, with some molecules remaining intact while others produce H^+ and the conjugate base.

This balance is what allows buffer solutions to effectively neutralize acids and bases. When an acid is added to a buffer, the equilibrium shifts slightly to offset the added ions, maintaining the pH. This balancing act is crucial in numerous biochemical and industrial processes where stable pH environments are needed.

Acid-Base Reactions

Acid-base reactions involve the transfer of H^+ ions from an acid to a base. These reactions are fundamental in chemistry, driving processes like neutralization, where an acid and base react to form water and a salt.

In the creation of buffer solutions, understanding acid-base reactions is essential. When a strong base is added to a weak acid, some of the acid molecules will donate H^+ ions to neutralize the base, forming water and leaving the conjugate base. This particular reaction is vital for forming buffer solutions as demonstrated in the exercise.

For instance, in option (c), HF partially reacts with NaOH , producing F^- and some remaining HF , establishing a buffer system. The reaction doesn’t go to completion, allowing for the coexistence of both the weak acid and the conjugate base, which keeps the solution's pH relatively steady even when additional acids or bases are introduced.