Question

Determine whether or not the mixing of each pair of solutions results in a buffer. a. \(75.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{HF} ; 55.0 \mathrm{~mL}\) of \(0.15 \mathrm{M} \mathrm{NaF}\) b. \(150.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{HF} ; 135.0 \mathrm{~mL}\) of \(0.175 \mathrm{M} \mathrm{HCl}\) c. \(165.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{HF} ; 135.0 \mathrm{~mL}\) of \(0.050 \mathrm{M} \mathrm{KOH}\) d. \(125.0 \mathrm{~mL}\) of \(0.15 \mathrm{M} \mathrm{CH}_{3} \mathrm{NH}_{2} ; 120.0 \mathrm{~mL}\) of \(0.25 \mathrm{M} \mathrm{CH}_{3} \mathrm{NH}_{3} \mathrm{Cl}\) e. \(105.0 \mathrm{~mL}\) of \(0.15 \mathrm{M} \mathrm{CH}_{3} \mathrm{NH}_{2} ; 95.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{HCl}\)

Short answer

Mixture (a) forms a buffer, (b) does not form a buffer, (c) does not form a buffer, (d) forms a buffer, and (e) forms a buffer.

Step-by-step solution

Understanding What Constitutes a Buffer Solution

A buffer solution consists of a weak acid and its conjugate base or a weak base and its conjugate acid. To form a buffer, the solution must contain significant amounts of both species so they can resist changes in pH when small amounts of acid or base are added.

Analyze Mixture (a)

HF is a weak acid, and NaF is a salt containing its conjugate base (F-). These components can create a buffer. Calculate the moles of each: HF (75.0 mL * 0.10 M = 7.5 mmol), NaF (55.0 mL * 0.15 M = 8.25 mmol). Both components are present in similar quantities, so mixture (a) can act as a buffer.

Analyze Mixture (b)

HF is a weak acid, but HCl is a strong acid, it does not have a conjugate base that can buffer the solution. Mixture (b) is not a buffer.

Analyze Mixture (c)

HF is a weak acid and KOH is a strong base. The strong base will completely neutralize some of the weak acid, possibly leaving excess weak acid but not its conjugate base. Without significant amounts of both, mixture (c) cannot act as a buffer.

Analyze Mixture (d)

CH3NH2 is a weak base and CH3NH3Cl is its conjugate acid (the chloride is a spectator ion). Calculate moles: CH3NH2 (125.0 mL * 0.15 M = 18.75 mmol), CH3NH3Cl (120.0 mL * 0.25 M = 30 mmol). Both components are present, so mixture (d) should act as a buffer.

Analyze Mixture (e)

CH3NH2 is a weak base and HCl is a strong acid. The strong acid will completely neutralize some of the weak base, creating its conjugate acid (CH3NH3+). Because we started with similar amounts of weak base and strong acid, mixture (e) should also act as a buffer.

Key concepts

Acid-Base Equilibria

Acid-base equilibria are fundamental to understanding buffer solutions. This crucial concept deals with the balance between acids and bases in a solution. An acid donates protons (H+ ions), while a base accepts them. The equilibrium is the state where the rate of the forward reaction (acid donating protons) equals the rate of the reverse reaction (base accepting protons).

In the context of buffer solutions, acid-base equilibrium maintains the pH of the solution despite the addition of small amounts of an acid or a base. This balancing act is possible because a buffer is comprised of a weak acid and its conjugate base, or a weak base and its conjugate acid, each partaking in reversible reactions to neutralize added acids or bases without significantly changing the pH.

To determine whether a mixture acts as a buffer, as in the given exercise, we must identify if both components of the buffer (the weak acid/base and its conjugate) are present and in significant amounts. Only then can the mixture resist shifts in pH, demonstrating the concept of acid-base equilibria in action.

Weak Acids and Bases

Weak acids and bases play a pivotal role in buffer solutions. A weak acid partially dissociates in water, releasing some protons to form its conjugate base, but not all acid molecules dissociate. This incomplete dissociation is key to buffer action. Similarly, a weak base partially accepts protons from water, forming its conjugate acid.

For instance, HF (hydrofluoric acid) is a weak acid that doesn't completely dissociate in solution. In the exercise, HF is used to possibly form a buffer. It's essential for students to grasp that only weak acids or bases, as opposed to strong ones, can form buffers because they retain both species (the acid/base and its conjugate) in the solution. If we were to use a strong acid or base, it would fully dissociate or accept protons, leaving no weak component to help maintain the pH equilibrium. The exercise clearly illustrates this difference, highlighting that solutions containing strong acids (like HCl) or strong bases (like KOH) do not result in a buffer.

Conjugate Acid-Base Pairs

The concept of conjugate acid-base pairs is central to understanding buffer chemistry. When a weak acid releases a proton, it becomes its conjugate base. Conversely, when a weak base accepts a proton, it transforms into its conjugate acid. These pairs are always present in equilibrium in a buffer solution.

Let's look at the mixture (d) from the exercise as an example. CH3NH2 is a weak base, and CH3NH3Cl is the salt that contains its conjugate acid (CH3NH3+). They work together to form a buffer. When an additional acid is added to such a mixture, the conjugate base (in this case, CH3NH2) accepts protons, preventing a significant pH change. If a base is added, the conjugate acid (CH3NH3+) donates protons.

Understanding conjugate acid-base pairs allows students to predict how a substance will behave in a solution and whether it can form part of a buffer system. This relationship also helps in identifying if a buffer exists when mixing solutions, as demonstrated in the exercise with multiple mixtures and their components.