Define buffer solution .

A solution that resists significant changes in pH when small amounts of acid or base are added. It consists of a weak acid and its conjugate base (acidic buffer) or a weak base and its conjugate acid (basic buffer).

True or False? A buffer solution can maintain a constant pH regardless of how much acid or base is added.

False. A buffer can only resist pH change when small amounts of acid or base are added. Adding large (excessive) amounts will exhaust the buffer components and cause a significant pH change.

How does an acidic buffer resist a rise in pH when a small amount of base (OH - ) is added?

The added OH - reacts with H + to form water, shifting the equilibrium CH 3 COOH ⇌ H + + CH 3 COO - to the right so that CH 3 COOH dissociates to replenish H + . Because both [CH 3 COOH] and [CH 3 COO - ] are large reserves, the ratio [CH 3 COO - ]/[CH 3 COOH] barely changes, so pH remains nearly constant.

True or False? An acidic buffer solution must contain a strong acid and its conjugate base.

False. An acidic buffer must contain a weak acid and its conjugate base. A strong acid fully dissociates and cannot re-establish equilibrium to replenish H + , so it cannot act as a buffer.

Why is a buffer most effective when the pH of the solution equals the pKa of the weak acid?

When pH = pK a , [HA] = [A - ] (equal concentrations of conjugate acid and conjugate base ). This means the buffer has the maximum capacity to absorb both added H + (neutralised by A - ) and added OH - (neutralised by HA) without a large shift in the [A - ]/[HA] ratio, giving the greatest resistance to pH change in either direction.

An acidic buffer contains a weak acid and its .......... , while a basic buffer contains a weak base and its .......... , both in relatively high concentrations to resist pH change.

An acidic buffer contains a weak acid and its conjugate base , while a basic buffer contains a weak base and its conjugate acid , both in relatively high concentrations to resist pH change.

Define the Henderson-Hasselbalch equation .

An equation used to calculate the pH of a buffer solution: pH = pK a + log([A - ]/[HA]) , where [A - ] is the molar concentration of the conjugate base and [HA] is the molar concentration of the weak acid.

True or False? When [A - ] > [HA] in a buffer, the pH is greater than the pKa.

True. If [A - ] > [HA], then log([A - ]/[HA]) > 0, so pH = pK a + (positive number) > pK a . A higher conjugate base concentration makes the solution less acidic.

Why does doubling both [HA] and [A - ] by the same factor leave the pH of a buffer unchanged?

The Henderson-Hasselbalch equation shows pH depends on the ratio [A - ]/[HA], not on the absolute concentrations. If both concentrations are multiplied by the same factor, the ratio is unchanged, log([A - ]/[HA]) is unchanged, and therefore pH = pK a + log([A - ]/[HA]) remains the same.

A buffer contains 0.400 M acetic acid and 0.200 M sodium acetate. The pK a of acetic acid is 4.76. How is the Henderson–Hasselbalch equation applied to find the pH?

pH = pK a + log([A - ] / [HA]) = 4.76 + log(0.200 / 0.400) = 4.76 + log(0.500) = 4.46 \n\nThe ratio [A - ]/[HA] < 1, so pH < pK a — consistent with the buffer having more acetic acid than conjugate base.

A buffer contains 0.305 M acetic acid and 0.520 M sodium acetate. The pKa of acetic acid is 4.76. What is the pH of this buffer, and how is the Henderson-Hasselbalch equation applied?

Using pH = pK a + log([A - ]/[HA]):\n\npH = 4.76 + log(0.520 / 0.305)\n\npH = 4.76 + log(1.705)\n\npH = 4.76 + 0.23 = 4.99 \n\nThe log term is positive because [A - ] > [HA], shifting the pH slightly above pK a .

The Henderson-Hasselbalch equation states that pH = pK a + log( .......... / .......... ), where the numerator is the conjugate base concentration and the denominator is the weak acid concentration.

The Henderson-Hasselbalch equation states that pH = pK a + log( [A - ] / [HA] ), where the numerator is the conjugate base concentration and the denominator is the weak acid concentration.

Define buffer capacity .

The amount of acid or base that can be added to a buffer before its pH changes significantly. Greater buffer capacity means the buffer can absorb more H + or OH - ions without a large pH shift.

True or False? A buffer containing 0.80 M acetic acid / 0.40 M acetate has greater capacity than one containing 0.080 M acetic acid / 0.040 M acetate.

True. Both buffers have the same initial pH (same ratio), but the higher-concentration buffer has larger reserve supplies of both CH 3 COOH and CH 3 COO - . It can absorb more moles of added acid or base before the ratio changes enough to cause a significant pH shift.

Why does a buffer with equal concentrations of weak acid and conjugate base have the highest buffer capacity?

When [HA] = [A - ], the buffer can absorb equal amounts of added H + (consumed by A - ) and added OH - (consumed by HA), so neither component is limiting and the resistance to pH change in both directions is maximised. Any imbalance in the ratio — e.g. [HA] >> [A - ] — limits the buffer capacity in one direction.

True or False? Two buffers with the same [HA]/[A - ] ratio always have the same buffer capacity, regardless of their absolute concentrations.

False. Buffer capacity depends on the absolute concentrations of HA and A⁻, not just their ratio. Two buffers with the same ratio but different concentrations will have the same pH, but the higher-concentration buffer has more moles of each component to neutralise added acid or base, giving it a greater capacity.

When H + ions are added to an acetic acid / sodium acetate buffer, what reaction occurs and how does this maintain pH?

The added H + reacts with the large reserve of CH 3 COO - : CH 3 COO - (aq) + H + (aq) → CH 3 COOH (aq). This consumes the H + before it can lower the pH significantly. Because [CH 3 COO - ] and [CH 3 COOH] are both large, the change in their ratio — and therefore the change in pH — is very small.

Buffer capacity increases when the .......... concentrations of the weak acid and conjugate base are increased, even if their .......... stays the same.

Buffer capacity increases when the absolute concentrations of the weak acid and conjugate base are increased, even if their ratio stays the same.

Define pH-dependent solubility .

The phenomenon where the solubility of a salt changes with pH because the ions produced on dissolving can react with H 3 O + or OH - in solution, shifting the dissolution equilibrium and altering how much solid dissolves.

True or False? Decreasing the pH (adding H 3 O + ) always increases the solubility of a sparingly soluble salt.

False. Decreasing pH increases solubility only if the salt produces a basic anion (weak conjugate base) that reacts with H 3 O + . If the anion is the conjugate base of a strong acid (e.g. NO 3 - , Cl - ), it is negligible and pH has no effect on solubility.

Why does lowering the pH increase the solubility of Fe(OH) 3 ?

In the equilibrium Fe(OH) 3 (s) ⇌ Fe 3+ (aq) + 3OH - (aq), adding H 3 O + consumes OH - ions (H 3 O + + OH - → 2H 2 O), and by Le Chatelier's principle this removes a product and shifts the equilibrium right, dissolving more Fe(OH) 3 and increasing solubility.

True or False? Adding OH - to a saturated solution of Fe(OH) 3 increases its solubility.

False. Adding OH - creates an excess of a product ion (common-ion effect). By Le Chatelier's principle, the equilibrium shifts left, causing more Fe(OH) 3 to precipitate. Solubility therefore decreases.

Why does changing pH affect the solubility of BaF 2 when acid is added, but not when base is added?

In BaF 2 (s) ⇌ Ba 2+ (aq) + 2F - (aq), adding acid removes the weak conjugate base F - (F - + H 3 O + → HF + H 2 O), shifting equilibrium right and increasing solubility. Ba 2+ does not react with OH - , so adding base has no effect on the equilibrium and solubility is unchanged.

The solubility of a salt is affected by pH only when its ions can react with .......... or .......... in solution; ions that cannot do this are described as negligible.

The solubility of a salt is affected by pH only when its ions can react with H 3 O + or OH - in solution; ions that cannot do this are described as negligible.

AP®ChemistryCollege BoardFlashcardsUnit 8: Acids & BasesBuffers

Buffers (College Board AP® Chemistry): Flashcards

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Define buffer solution.

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Define buffer solution.
A solution that resists significant changes in pH when small amounts of acid or base are added. It consists of a weak acid and its conjugate base (acidic buffer) or a weak base and its conjugate acid (basic buffer).
True or False?
A buffer solution can maintain a constant pH regardless of how much acid or base is added.
False.
A buffer can only resist pH change when small amounts of acid or base are added. Adding large (excessive) amounts will exhaust the buffer components and cause a significant pH change.
How does an acidic buffer resist a rise in pH when a small amount of base (OH^-) is added?
The added OH^- reacts with H⁺ to form water, shifting the equilibrium CH₃COOH ⇌ H⁺ + CH₃COO^- to the right so that CH₃COOH dissociates to replenish H⁺. Because both [CH₃COOH] and [CH₃COO^-] are large reserves, the ratio [CH₃COO^-]/[CH₃COOH] barely changes, so pH remains nearly constant.
True or False?
An acidic buffer solution must contain a strong acid and its conjugate base.
False.
An acidic buffer must contain a weak acid and its conjugate base. A strong acid fully dissociates and cannot re-establish equilibrium to replenish H⁺, so it cannot act as a buffer.
Why is a buffer most effective when the pH of the solution equals the pKa of the weak acid?
When pH = pK_a, [HA] = [A^-] (equal concentrations of conjugate acid and conjugate base). This means the buffer has the maximum capacity to absorb both added H⁺ (neutralised by A^-) and added OH^- (neutralised by HA) without a large shift in the [A^-]/[HA] ratio, giving the greatest resistance to pH change in either direction.
An acidic buffer contains a weak acid and its .........., while a basic buffer contains a weak base and its .........., both in relatively high concentrations to resist pH change.
An acidic buffer contains a weak acid and its conjugate base, while a basic buffer contains a weak base and its conjugate acid, both in relatively high concentrations to resist pH change.
Define the Henderson-Hasselbalch equation.
An equation used to calculate the pH of a buffer solution: pH = pK_a + log([A^-]/[HA]), where [A^-] is the molar concentration of the conjugate base and [HA] is the molar concentration of the weak acid.
True or False?
When [A^-] > [HA] in a buffer, the pH is greater than the pKa.
True.
If [A^-] > [HA], then log([A^-]/[HA]) > 0, so pH = pK_a + (positive number) > pK_a. A higher conjugate base concentration makes the solution less acidic.
Why does doubling both [HA] and [A^-] by the same factor leave the pH of a buffer unchanged?
The Henderson-Hasselbalch equation shows pH depends on the ratio [A^-]/[HA], not on the absolute concentrations. If both concentrations are multiplied by the same factor, the ratio is unchanged, log([A^-]/[HA]) is unchanged, and therefore pH = pK_a + log([A^-]/[HA]) remains the same.
A buffer contains 0.400 M acetic acid and 0.200 M sodium acetate. The pK_a of acetic acid is 4.76. How is the Henderson–Hasselbalch equation applied to find the pH?
pH = pK_a + log([A^-] / [HA]) = 4.76 + log(0.200 / 0.400) = 4.76 + log(0.500) = 4.46\n\nThe ratio [A^-]/[HA] < 1, so pH < pK_a — consistent with the buffer having more acetic acid than conjugate base.
A buffer contains 0.305 M acetic acid and 0.520 M sodium acetate. The pKa of acetic acid is 4.76. What is the pH of this buffer, and how is the Henderson-Hasselbalch equation applied?
Using pH = pK_a + log([A^-]/[HA]):\n\npH = 4.76 + log(0.520 / 0.305)\n\npH = 4.76 + log(1.705)\n\npH = 4.76 + 0.23 = 4.99\n\nThe log term is positive because [A^-] > [HA], shifting the pH slightly above pK_a.
The Henderson-Hasselbalch equation states that pH = pK_a + log(.......... / .......... ), where the numerator is the conjugate base concentration and the denominator is the weak acid concentration.
The Henderson-Hasselbalch equation states that pH = pK_a + log([A^-] / [HA] ), where the numerator is the conjugate base concentration and the denominator is the weak acid concentration.
Define buffer capacity.
The amount of acid or base that can be added to a buffer before its pH changes significantly. Greater buffer capacity means the buffer can absorb more H⁺ or OH^- ions without a large pH shift.
True or False?
A buffer containing 0.80 M acetic acid / 0.40 M acetate has greater capacity than one containing 0.080 M acetic acid / 0.040 M acetate.
True.
Both buffers have the same initial pH (same ratio), but the higher-concentration buffer has larger reserve supplies of both CH₃COOH and CH₃COO^-. It can absorb more moles of added acid or base before the ratio changes enough to cause a significant pH shift.
Why does a buffer with equal concentrations of weak acid and conjugate base have the highest buffer capacity?
When [HA] = [A^-], the buffer can absorb equal amounts of added H⁺ (consumed by A^-) and added OH^- (consumed by HA), so neither component is limiting and the resistance to pH change in both directions is maximised. Any imbalance in the ratio — e.g. [HA] >> [A^-] — limits the buffer capacity in one direction.
True or False?
Two buffers with the same [HA]/[A^-] ratio always have the same buffer capacity, regardless of their absolute concentrations.
False.
Buffer capacity depends on the absolute concentrations of HA and A⁻, not just their ratio. Two buffers with the same ratio but different concentrations will have the same pH, but the higher-concentration buffer has more moles of each component to neutralise added acid or base, giving it a greater capacity.
When H⁺ ions are added to an acetic acid / sodium acetate buffer, what reaction occurs and how does this maintain pH?
The added H⁺ reacts with the large reserve of CH₃COO^-: CH₃COO^- (aq) + H⁺ (aq) → CH₃COOH (aq). This consumes the H⁺ before it can lower the pH significantly. Because [CH₃COO^-] and [CH₃COOH] are both large, the change in their ratio — and therefore the change in pH — is very small.
Buffer capacity increases when the .......... concentrations of the weak acid and conjugate base are increased, even if their .......... stays the same.
Buffer capacity increases when the absolute concentrations of the weak acid and conjugate base are increased, even if their ratio stays the same.
Define pH-dependent solubility.
The phenomenon where the solubility of a salt changes with pH because the ions produced on dissolving can react with H₃O⁺ or OH^- in solution, shifting the dissolution equilibrium and altering how much solid dissolves.
True or False?
Decreasing the pH (adding H₃O⁺) always increases the solubility of a sparingly soluble salt.
False.
Decreasing pH increases solubility only if the salt produces a basic anion (weak conjugate base) that reacts with H₃O⁺. If the anion is the conjugate base of a strong acid (e.g. NO₃^-, Cl^-), it is negligible and pH has no effect on solubility.
Why does lowering the pH increase the solubility of Fe(OH)₃?
In the equilibrium Fe(OH)₃ (s) ⇌ Fe³⁺ (aq) + 3OH^- (aq), adding H₃O⁺ consumes OH^- ions (H₃O⁺ + OH^- → 2H₂O), and by Le Chatelier's principle this removes a product and shifts the equilibrium right, dissolving more Fe(OH)₃ and increasing solubility.
True or False?
Adding OH^- to a saturated solution of Fe(OH)₃ increases its solubility.
False.
Adding OH^- creates an excess of a product ion (common-ion effect). By Le Chatelier's principle, the equilibrium shifts left, causing more Fe(OH)₃ to precipitate. Solubility therefore decreases.
Why does changing pH affect the solubility of BaF₂ when acid is added, but not when base is added?
In BaF₂ (s) ⇌ Ba²⁺ (aq) + 2F^- (aq), adding acid removes the weak conjugate base F^- (F^- + H₃O⁺ → HF + H₂O), shifting equilibrium right and increasing solubility. Ba²⁺ does not react with OH^-, so adding base has no effect on the equilibrium and solubility is unchanged.
The solubility of a salt is affected by pH only when its ions can react with .......... or .......... in solution; ions that cannot do this are described as negligible.
The solubility of a salt is affected by pH only when its ions can react with H₃O⁺ or OH^- in solution; ions that cannot do this are described as negligible.

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