Acid-Base Reactions and Buffers - AP Chemistry

Two buffer solutions are prepared: Buffer X contains a weak acid $\mathrm{HA}$ and its conjugate base $\mathrm{A^-}$ , with $\mathrm{HA}$ in excess. Buffer Y contains the same conjugate pair, but with $\mathrm{A^-}$ in excess. A small amount of $\mathrm{NaOH(aq)}$ is added to each buffer. Which statement best describes the buffering reaction that occurs in both solutions?
Answer: In both buffers, $\mathrm{HA}$ is the primary species that reacts with added $\mathrm{OH^-}$ to form $\mathrm{A^-}$ and $\mathrm{H_2O}$ . This question tests understanding of properties of buffers. Both Buffer X and Buffer Y contain the same weak acid HA and conjugate base A⁻ pair, just in different ratios, allowing both to resist pH changes when small amounts of acid or base are added. When NaOH is added to either buffer, the OH⁻ ions react primarily with the weak acid HA to form A⁻ and H₂O: HA + OH⁻ → A⁻ + H₂O. This reaction occurs in both buffers regardless of which component is in excess, as long as sufficient HA is present to react with the added base. Choice A is incorrect because A⁻ is the conjugate base and cannot react with OH⁻ to form HA—this would require a base to react with another base to form an acid, which violates acid-base principles. When analyzing buffer reactions, remember that the weak acid component always reacts with added base, regardless of the initial ratio of buffer components.
A buffer solution contains $\text{HNO}_2(aq)$ and $\text{NO}_2^-(aq)$ in comparable amounts. A small amount of $\text{NaOH}(aq)$ is added. Which statement best describes the primary acid–base reaction that occurs?
Answer: $\text{HNO}_2$ reacts with $\text{OH}^-$ to form $\text{NO}_2^-$ and $\text{H}_2\text{O}$ , limiting the pH increase. This question assesses the properties of buffers. Buffers contain a weak acid and its conjugate base, which work together to resist pH changes. When base is added, the weak acid reacts with the added OH⁻ to form more conjugate base and water, consuming the OH⁻ and preventing a large rise in pH. When acid is added, the conjugate base reacts with the added H⁺ to form more weak acid, consuming the H⁺ and preventing a large drop in pH. A common misconception is that buffers keep pH constant by preventing reactions, as in choice E, but actually, they rely on reactions between components and added species. To solve buffer problems, identify which buffer component reacts with the added species—the weak acid neutralizes added base, and the conjugate base neutralizes added acid.
A buffer is made by combining $\mathrm{H_2PO_4^-(aq)}$ and $\mathrm{HPO_4^{2-}(aq)}$ so that both species are present in comparable amounts. A small amount of $\mathrm{HCl(aq)}$ is added. Which change occurs as the buffer responds?
Answer: The concentration of $\mathrm{HPO_4^{2-}}$ decreases as it is converted to $\mathrm{H_2PO_4^-}$ . This question tests understanding of properties of buffers. The phosphate buffer contains H₂PO₄⁻ (weak acid) and HPO₄²⁻ (conjugate base) in comparable amounts, enabling it to resist pH changes when small amounts of acid or base are added. When HCl is added, it provides H⁺ ions that react primarily with the conjugate base HPO₄²⁻ to form H₂PO₄⁻: HPO₄²⁻ + H⁺ → H₂PO₄⁻. This reaction causes the concentration of HPO₄²⁻ to decrease while the concentration of H₂PO₄⁻ increases, consuming most of the added acid and preventing a large pH decrease. Choice C is incorrect because Cl⁻ is the conjugate base of a strong acid (HCl) and is too weak a base to react with H⁺ to form HClO under these conditions. When predicting buffer responses, remember that conjugate bases react with added acids, causing their concentration to decrease as they are converted to the weak acid form.
A buffer is made using carbonic acid, $\mathrm{H_2CO_3(aq)}$ , and bicarbonate, $\mathrm{HCO_3^-(aq)}$ , with approximately equal amounts of each present. A small amount of $\mathrm{NaOH(aq)}$ is added. Which species is primarily consumed in the reaction that limits the pH change?
Answer: $\mathrm{H_2CO_3}$ This question tests understanding of properties of buffers. The carbonic acid/bicarbonate buffer contains H₂CO₃ (weak acid) and HCO₃⁻ (conjugate base) in approximately equal amounts, allowing it to resist pH changes when small amounts of acid or base are added. When NaOH is added, it provides OH⁻ ions that react primarily with the weak acid H₂CO₃ to form HCO₃⁻ and H₂O: H₂CO₃ + OH⁻ → HCO₃⁻ + H₂O. This reaction consumes the added base and H₂CO₃, preventing a large increase in pH. Choice E is incorrect because NaOH is a strong base that provides OH⁻ ions, not H₃O⁺ ions—this represents a fundamental misunderstanding of acid-base chemistry. To identify the consumed species in a buffer reaction, determine which buffer component can neutralize the added species: weak acids neutralize added bases.
A buffer is prepared by mixing $\mathrm{CH_3NH_2(aq)}$ (a weak base) and $\mathrm{CH_3NH_3Cl(aq)}$ so that both $\mathrm{CH_3NH_2}$ and $\mathrm{CH_3NH_3^+}$ are present. A small amount of strong acid is added. Which statement best explains why the pH does not drop as much as it would in pure water?
Answer: The added $\mathrm{H^+}$ is primarily consumed by $\mathrm{CH_3NH_2}$ to form $\mathrm{CH_3NH_3^+}$ , reducing the amount of free $\mathrm{H^+}$ . This question tests understanding of properties of buffers. The methylamine buffer contains CH₃NH₂ (weak base) and CH₃NH₃⁺ (conjugate acid), allowing it to resist pH changes when small amounts of acid or base are added. When strong acid is added, it provides H⁺ ions that react primarily with the weak base CH₃NH₂ to form CH₃NH₃⁺: CH₃NH₂ + H⁺ → CH₃NH₃⁺. This reaction consumes most of the added H⁺, preventing the large pH decrease that would occur in pure water where all H⁺ remains free in solution. Choice D is incorrect because CH₃NH₂ is a weak base, not a strong base, and buffers do not completely neutralize all added acid—they only minimize pH changes. To understand buffer action, identify which component neutralizes the added species: weak bases consume added acids, while conjugate acids consume added bases.
A buffer is made from the weak acid $\text{HF}(aq)$ and its conjugate base $\text{F}^-(aq)$ , with an excess of $\text{F}^-$ present. A small amount of strong acid is added. Which statement best describes the immediate chemical reaction that helps resist the pH change?
Answer: $\text{F}^-$ reacts with added $\text{H}^+$ to form $\text{HF}$ , reducing the increase in $\text{H}_3\text{O}^+$ concentration. This question assesses the properties of buffers. Buffers contain a weak acid and its conjugate base, which work together to resist pH changes. When acid is added, the conjugate base reacts with the added H⁺ to form more weak acid, consuming the H⁺ and preventing a large drop in pH. When base is added, the weak acid reacts with the added OH⁻ to form more conjugate base and water, consuming the OH⁻ and preventing a large rise in pH. A common misconception is that buffers prevent acids from dissociating entirely, as in choice E, but actually, buffers react with added species to minimize pH shifts rather than stopping dissociation. To solve buffer problems, identify which buffer component reacts with the added species—the conjugate base neutralizes added acid, and the weak acid neutralizes added base.
A buffer is prepared using $\text{C}_6\text{H}_5\text{COOH}(aq)$ (benzoic acid) and $\text{C}_6\text{H}_5\text{COO}^-(aq)$ in comparable amounts. A small amount of $\text{NaOH}(aq)$ is added. Which statement best describes what happens to the pH and why?
Answer: The pH increases only slightly because $\text{C}_6\text{H}_5\text{COOH}$ neutralizes added $\text{OH}^-$ to form $\text{C}_6\text{H}_5\text{COO}^-$ and water. This question assesses the properties of buffers. Buffers contain a weak acid and its conjugate base, which work together to resist pH changes. When base is added, the weak acid reacts with the added OH⁻ to form more conjugate base and water, consuming the OH⁻ and preventing a large rise in pH. When acid is added, the conjugate base reacts with the added H⁺ to form more weak acid, consuming the H⁺ and preventing a large drop in pH. A common misconception is that buffers keep pH exactly constant regardless of additions, as in choice D, but small changes do occur. To solve buffer problems, identify which buffer component reacts with the added species—the weak acid neutralizes added base, and the conjugate base neutralizes added acid.
A student prepares a buffer by mixing aqueous acetic acid, $\mathrm{HC_2H_3O_2}$ , and sodium acetate, $\mathrm{NaC_2H_3O_2}$ , so that the solution contains comparable amounts of $\mathrm{HC_2H_3O_2}$ and $\mathrm{C_2H_3O_2^-}$ . The student then adds a small amount of $\mathrm{HCl(aq)}$ to the buffer. Which statement best describes what happens in the solution?
Answer: The added $\mathrm{H^+}$ is primarily consumed by $\mathrm{C_2H_3O_2^-}$ to form $\mathrm{HC_2H_3O_2}$ , so the pH decreases only slightly. This question tests understanding of properties of buffers. A buffer solution contains both a weak acid (HC₂H₃O₂) and its conjugate base (C₂H₃O₂⁻), which allows it to resist pH changes when small amounts of acid or base are added. When HCl is added, the H⁺ ions from the strong acid are consumed by the acetate ions (C₂H₃O₂⁻) through the reaction: C₂H₃O₂⁻ + H⁺ → HC₂H₃O₂. This neutralization reaction prevents most of the added H⁺ from remaining free in solution, which would otherwise cause a sharp pH decrease. Option D incorrectly suggests that buffers only neutralize bases, not acids—this is a common misconception since buffers work bidirectionally. The key strategy is to identify which buffer component (the base form) reacts with added acid to minimize pH change.
A buffer contains carbonic acid, $\text{H}_2\text{CO}_3(aq)$ , and bicarbonate, $\text{HCO}_3^-(aq)$ , in comparable amounts. A small amount of $\text{HCl}(aq)$ is added. Which change best describes what happens to the relative amounts of buffer components?
Answer: $\text{HCO}_3^-$ decreases and $\text{H}_2\text{CO}_3$ increases because $\text{HCO}_3^-$ reacts with added $\text{H}^+$ . This question assesses the properties of buffers. Buffers contain a weak acid and its conjugate base, which work together to resist pH changes. When acid is added, the conjugate base reacts with the added H⁺ to form more weak acid, consuming the H⁺ and preventing a large drop in pH. When base is added, the weak acid reacts with the added OH⁻ to form more conjugate base and water, consuming the OH⁻ and preventing a large rise in pH. A common misconception is that buffer concentrations remain completely unchanged, as in choice E, but actually, the relative amounts shift slightly as one component reacts. To solve buffer problems, identify which buffer component reacts with the added species—the conjugate base neutralizes added acid, and the weak acid neutralizes added base.
A student prepares a buffer by mixing acetic acid, $\mathrm{HC_2H_3O_2}$ , and sodium acetate, $\mathrm{NaC_2H_3O_2}$ , so that the solution contains comparable amounts of $\mathrm{HC_2H_3O_2(aq)}$ and $\mathrm{C_2H_3O_2^-(aq)}$ . A small amount of $\mathrm{HCl(aq)}$ is added. Which statement best describes the primary reaction that helps the solution resist a large change in pH?
Answer: $\mathrm{C_2H_3O_2^-}$ reacts with added $\mathrm{H^+}$ to form $\mathrm{HC_2H_3O_2}$ , removing most of the added $\mathrm{H^+}$ . This question tests understanding of properties of buffers. A buffer contains a weak acid (HC₂H₃O₂) and its conjugate base (C₂H₃O₂⁻) in comparable amounts, which allows it to resist pH changes when small amounts of acid or base are added. When HCl is added, it provides H⁺ ions that are primarily consumed by the conjugate base C₂H₃O₂⁻, forming more of the weak acid HC₂H₃O₂ according to the reaction: C₂H₃O₂⁻ + H⁺ → HC₂H₃O₂. This reaction removes most of the added H⁺ from solution, preventing a large decrease in pH. Choice A is incorrect because HC₂H₃O₂ is already a weak acid and cannot accept another proton to form H₂C₂H₃O₂⁺ under normal conditions. To identify how a buffer responds to added acid or base, determine which buffer component (weak acid or conjugate base) can react with the added species—the conjugate base reacts with added acid, while the weak acid reacts with added base.
A buffer contains $\text{H}_3\text{C}_6\text{H}_5\text{O}_7(aq)$ (citric acid, treated here as a weak acid $\text{HA}$ ) and its conjugate base $\text{A}^-(aq)$ in comparable amounts. A small amount of strong acid is added. Which statement best predicts how the concentration of $\text{A}^-$ changes?
Answer: $[\text{A}^-]$ decreases because $\text{A}^-$ reacts with added $\text{H}^+$ to form $\text{HA}$ . This question assesses the properties of buffers. Buffers contain a weak acid and its conjugate base, which work together to resist pH changes. When acid is added, the conjugate base reacts with the added H⁺ to form more weak acid, consuming the H⁺ and preventing a large drop in pH. When base is added, the weak acid reacts with the added OH⁻ to form more conjugate base and water, consuming the OH⁻ and preventing a large rise in pH. A common misconception is that conjugate base concentrations stay fixed, as in choice C, but they decrease when reacting with added acid. To solve buffer problems, identify which buffer component reacts with the added species—the conjugate base neutralizes added acid, and the weak acid neutralizes added base.
A buffer contains $\text{H}_2\text{PO}_4^-(aq)$ and $\text{HPO}_4^{2-}(aq)$ in comparable amounts. A small amount of $\text{HCl}(aq)$ is added. Which statement best describes the main reason the pH change is small?
Answer: The added $\text{H}^+$ is consumed by $\text{HPO}_4^{2-}$ to form $\text{H}_2\text{PO}_4^-$ , so little free $\text{H}_3\text{O}^+$ remains. This question assesses the properties of buffers. Buffers contain a weak acid and its conjugate base, which work together to resist pH changes. When acid is added, the conjugate base reacts with the added H⁺ to form more weak acid, consuming the H⁺ and preventing a large drop in pH. When base is added, the weak acid reacts with the added OH⁻ to form more conjugate base and water, consuming the OH⁻ and preventing a large rise in pH. A common misconception is that added acid does not dissociate in buffers, as in choice C, but it does, and the buffer reacts to consume the H⁺. To solve buffer problems, identify which buffer component reacts with the added species—the conjugate base neutralizes added acid, and the weak acid neutralizes added base.
A buffer is prepared by mixing a weak acid $\mathrm{HA}$ with its conjugate base $\mathrm{A^-}$ . A student mistakenly claims, "If you add a small amount of strong base, the pH will not change at all because buffers keep pH constant." Which statement best evaluates the claim?
Answer: The claim is incorrect; added $\mathrm{OH^-}$ reacts with $\mathrm{HA}$ to form $\mathrm{A^-}$ , so pH changes slightly rather than not at all. This question tests understanding of properties of buffers. The student's claim that pH will not change at all is incorrect because buffers minimize but do not completely eliminate pH changes. When a strong base like NaOH is added to an HA/A⁻ buffer, the OH⁻ ions react with the weak acid component: HA + OH⁻ → A⁻ + H₂O. This reaction consumes most of the added base and converts HA to A⁻, causing the ratio of conjugate base to weak acid to increase, which results in a slight pH increase. Option B incorrectly states that buffers contain no species that can react with OH⁻—the weak acid component specifically serves this purpose. The key concept is that buffers resist but don't prevent pH changes, and the strategy is to identify which component (weak acid) neutralizes added base.
A buffer is prepared with the weak base $\text{CH}_3\text{NH}_2(aq)$ and its conjugate acid $\text{CH}_3\text{NH}_3^+(aq)$ , with an excess of $\text{CH}_3\text{NH}_3^+$ . A small amount of strong base is added. Which statement best predicts the pH change?
Answer: The pH increases only slightly because $\text{CH}_3\text{NH}_3^+$ neutralizes added $\text{OH}^-$ to form $\text{CH}_3\text{NH}_2$ and $\text{H}_2\text{O}$ . This question assesses the properties of buffers. Buffers contain a weak base and its conjugate acid, which work together to resist pH changes. When base is added, the conjugate acid reacts with the added OH⁻ to form more weak base and water, consuming the OH⁻ and preventing a large rise in pH. When acid is added, the weak base reacts with the added H⁺ to form more conjugate acid, consuming the H⁺ and preventing a large drop in pH. A common misconception is that buffers eliminate added species completely without any pH change, as in choice C, but they only minimize the change, not prevent it entirely. To solve buffer problems, identify which buffer component reacts with the added species—the conjugate acid neutralizes added base, and the weak base neutralizes added acid.
Two solutions are prepared: - Solution 1: $\text{HCOOH}(aq)$ only (a weak acid) - Solution 2: a buffer made with $\text{HCOOH}(aq)$ and $\text{HCOO}^-(aq)$ in comparable amounts Equal small amounts of $\text{HCl}(aq)$ are added to each. Which statement best compares the pH changes?
Answer: Solution 2 shows a smaller pH decrease because $\text{HCOO}^-$ consumes much of the added $\text{H}^+$ . This question assesses the properties of buffers. Buffers contain a weak acid and its conjugate base, which work together to resist pH changes. When acid is added, the conjugate base reacts with the added H⁺ to form more weak acid, consuming the H⁺ and preventing a large drop in pH. When base is added, the weak acid reacts with the added OH⁻ to form more conjugate base and water, consuming the OH⁻ and preventing a large rise in pH. A common misconception is that buffers and weak acids behave the same way, but as shown here, buffers resist pH changes more effectively than weak acids alone, making choice A incorrect. To solve buffer problems, identify which buffer component reacts with the added species—the conjugate base neutralizes added acid, and the weak acid neutralizes added base.
A buffer solution contains nitrous acid, $\mathrm{HNO_2(aq)}$ , and sodium nitrite, $\mathrm{NaNO_2(aq)}$ , with $\mathrm{HNO_2}$ present in excess. A small amount of strong base is added. Which statement best describes the immediate effect on the buffer components?
Answer: $\mathrm{HNO_2}$ decreases as it reacts with $\mathrm{OH^-}$ to form $\mathrm{NO_2^-}$ and $\mathrm{H_2O}$ . This question tests understanding of properties of buffers. The nitrous acid/nitrite buffer contains HNO₂ (weak acid) and NO₂⁻ (conjugate base), with HNO₂ in excess, enabling it to resist pH changes when small amounts of acid or base are added. When strong base is added, it provides OH⁻ ions that react with the weak acid HNO₂ to form NO₂⁻ and H₂O: HNO₂ + OH⁻ → NO₂⁻ + H₂O. This reaction causes the concentration of HNO₂ to decrease while the concentration of NO₂⁻ increases, consuming most of the added base and preventing a large pH increase. Choice A is incorrect because NO₂⁻ is the conjugate base and cannot react with OH⁻ to form HNO₂—this would require the base to act as an acid, which is chemically impossible. When analyzing buffer component changes, remember that added base consumes the weak acid component, converting it to more conjugate base.
A student has two beakers: - Beaker 1 contains a buffer made from $\mathrm{CH_3COOH}$ and $\mathrm{CH_3COO^-}$ . - Beaker 2 contains pure water. The same small amount of $\mathrm{HCl(aq)}$ is added to each beaker and mixed. Which statement best compares the pH changes?
Answer: Beaker 1 experiences a smaller pH decrease because $\mathrm{CH_3COO^-}$ consumes much of the added $\mathrm{H^+}$ . This question tests understanding of properties of buffers. Beaker 1 contains a buffer with both CH₃COOH (weak acid) and CH₃COO⁻ (conjugate base), while Beaker 2 contains only pure water with no buffering capacity. When HCl is added to the buffer in Beaker 1, the H⁺ ions are largely consumed by the acetate ions: CH₃COO⁻ + H⁺ → CH₃COOH, resulting in only a small pH decrease. In contrast, when HCl is added to pure water in Beaker 2, all the H⁺ remains free in solution, causing a much larger pH decrease. Option C incorrectly suggests that buffers amplify pH changes—this misunderstands that buffers resist, not enhance, pH changes through neutralization reactions. The key insight is that buffers minimize pH changes by providing species that can neutralize added acids or bases, unlike pure water.
A student has a buffer made of $\text{CH}_3\text{COOH}(aq)$ and $\text{CH}_3\text{COO}^-(aq)$ . The student adds a small amount of $\text{CH}_3\text{COONa}(aq)$ (sodium acetate) without adding any strong acid or base. Which statement best predicts the immediate effect on pH?
Answer: The pH increases because adding $\text{CH}_3\text{COO}^-$ increases the solution's ability to consume $\text{H}^+$ , shifting the acid–base equilibrium toward fewer hydronium ions. This question assesses the properties of buffers. Buffers contain a weak acid and its conjugate base, which work together to resist pH changes. When acid is added, the conjugate base reacts with the added H⁺ to form more weak acid, consuming the H⁺ and preventing a large drop in pH. When base is added, the weak acid reacts with the added OH⁻ to form more conjugate base and water, consuming the OH⁻ and preventing a large rise in pH. A common misconception is that adding more conjugate base decreases pH, as in choice A, but it actually increases pH by shifting the equilibrium. To solve buffer problems, identify which buffer component reacts with the added species or how changes in ratio affect pH using the Henderson-Hasselbalch equation.
A student prepares a buffer by mixing equal amounts of hydrofluoric acid, $\mathrm{HF(aq)}$ , and sodium fluoride, $\mathrm{NaF(aq)}$ . The student then adds a small amount of $\mathrm{NaOH(aq)}$ . Which statement best describes what happens?
Answer: The added $\mathrm{OH^-}$ reacts primarily with $\mathrm{HF}$ to form $\mathrm{F^-}$ and $\mathrm{H_2O}$ , so the pH increases only slightly. This question tests understanding of properties of buffers. The buffer contains equal amounts of a weak acid (HF) and its conjugate base (F⁻), which enables the system to resist pH changes when small amounts of acid or base are added. When NaOH is added, the OH⁻ ions react with the weak acid HF to form F⁻ and water through the reaction: HF + OH⁻ → F⁻ + H₂O. This neutralization reaction consumes most of the added OH⁻, preventing a large increase in [OH⁻] and limiting the pH increase to only a slight change. Choice E is incorrect because it claims buffers behave the same as pure water when base is added—buffers specifically resist pH changes much more effectively than pure water due to the neutralization reactions. To analyze buffer behavior, identify which buffer component reacts with the added species: weak acids neutralize added bases, while conjugate bases neutralize added acids.
A buffer solution contains a weak acid $\mathrm{HA}$ and its conjugate base $\mathrm{A^-}$ , with $\mathrm{A^-}$ present in excess. A small amount of $\mathrm{HCl(aq)}$ is added. Which statement best describes the result?
Answer: The pH decreases only slightly because most added $\mathrm{H^+}$ is consumed by $\mathrm{A^-}$ to form $\mathrm{HA}$ . This question tests understanding of properties of buffers. The buffer contains a weak acid (HA) and its conjugate base (A⁻), with excess A⁻ present, enabling it to resist pH changes. When HCl is added, the H⁺ ions are consumed by the conjugate base (A⁻) through the reaction: A⁻ + H⁺ → HA. This neutralization prevents most of the added H⁺ from remaining free in solution, resulting in only a slight pH decrease rather than the sharp drop that would occur without the buffer. Option D incorrectly claims that buffers completely prevent any pH change—buffers minimize but don't eliminate pH changes, as the ratio of conjugate base to weak acid does shift slightly. The key strategy is recognizing that conjugate bases in buffers neutralize added acids, converting them to the weak acid form.

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