Acid-Base Reactions and Buffers
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AP Chemistry › Acid-Base Reactions and Buffers
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?
Both solutions show the same pH decrease because the same amount of $\text{HCl}$ is added.
Solution 1 shows a smaller pH decrease because weak acids resist pH change better than buffers.
Solution 2 shows a larger pH decrease because buffers contain more total solute.
Neither solution changes pH because $\text{HCl}$ is a strong acid and sets the pH.
Solution 2 shows a smaller pH decrease because $\text{HCOO}^-$ consumes much of the added $\text{H}^+$.
Explanation
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 student prepares a buffer by mixing equal concentrations of $\mathrm{HF(aq)}$ and $\mathrm{F^-(aq)}$ (from $\mathrm{NaF}$). The student then adds a small amount of $\mathrm{NaOH(aq)}$. Which statement best describes what happens?
The pH increases sharply because $\mathrm{HF}$ is a weak acid and cannot react with added $\mathrm{OH^-}$.
The added $\mathrm{OH^-}$ is primarily neutralized by $\mathrm{HF}$ to form $\mathrm{F^-}$ and $\mathrm{H_2O}$, so the pH increases only slightly.
The added $\mathrm{OH^-}$ reacts mainly with $\mathrm{Na^+}$, preventing a pH change.
The added $\mathrm{OH^-}$ is primarily neutralized by $\mathrm{F^-}$ to form $\mathrm{HF}$, so the pH decreases only slightly.
The pH remains exactly constant because the buffer converts all added $\mathrm{OH^-}$ into neutral salt with no equilibrium shift.
Explanation
This question tests understanding of properties of buffers. The HF/F⁻ buffer contains both a weak acid (HF) and its conjugate base (F⁻), allowing it to resist pH changes when acids or bases are added. When NaOH is added, the OH⁻ ions are neutralized by the weak acid component (HF) according to: HF + OH⁻ → F⁻ + H₂O. This reaction consumes most of the added hydroxide ions, preventing a sharp pH increase, though the pH does increase slightly as more F⁻ is formed and the ratio of F⁻ to HF increases. Option B incorrectly suggests that F⁻ reacts with OH⁻ to form HF—this is impossible as both F⁻ and OH⁻ are bases and cannot react in this way. The key strategy is to identify that weak acids in buffers neutralize added bases, while conjugate bases neutralize added acids.
A buffer contains the weak acid $\mathrm{H_2PO_4^-}$ and its conjugate base $\mathrm{HPO_4^{2-}}$ in comparable amounts. A small amount of $\mathrm{HCl(aq)}$ is added. Which reaction best represents the primary buffering process?
$\mathrm{H_2PO_4^- + H^+ \rightarrow H_3PO_4}$
$\mathrm{HPO_4^{2-} + H^+ \rightarrow H_2PO_4^-}$
$\mathrm{H_2PO_4^- + OH^- \rightarrow HPO_4^{2-} + H_2O}$
$\mathrm{HPO_4^{2-} + Cl^- \rightarrow HCl + PO_4^{3-}}$
$\mathrm{H^+ + OH^- \rightarrow H_2O}$ (from water only)
Explanation
This question tests understanding of properties of buffers. The phosphate buffer contains H₂PO₄⁻ (weak acid) and HPO₄²⁻ (conjugate base), which work together to resist pH changes. When HCl is added, the H⁺ ions are consumed by the conjugate base component (HPO₄²⁻) through the reaction: HPO₄²⁻ + H⁺ → H₂PO₄⁻. This neutralization converts the added strong acid into the weak acid form, preventing a sharp pH decrease and maintaining the buffer's effectiveness. Option C shows H₂PO₄⁻ accepting another proton, but this is not the primary buffering reaction since HPO₄²⁻ is more basic and reacts preferentially with added H⁺. The strategy for identifying buffer reactions is to recognize that the more basic component (higher charge on phosphate) neutralizes added acid.
A buffer solution is prepared with excess $\mathrm{H_2CO_3}$ and a smaller amount of $\mathrm{HCO_3^-}$ (from $\mathrm{NaHCO_3}$). A small amount of $\mathrm{HCl(aq)}$ is added. Which statement best describes the effect on the buffer?
The pH decreases sharply because the buffer contains excess weak acid and cannot react with added $\mathrm{H^+}$.
The added $\mathrm{H^+}$ reacts mainly with $\mathrm{Na^+}$, so the pH changes very little.
The added $\mathrm{H^+}$ is primarily consumed by $\mathrm{HCO_3^-}$ to form $\mathrm{H_2CO_3}$, so the pH decreases slightly.
The pH increases because adding $\mathrm{HCl}$ shifts $\mathrm{H_2CO_3}$ to produce more $\mathrm{HCO_3^-}$.
The pH remains exactly constant because buffers completely eliminate added $\mathrm{H^+}$.
Explanation
This question tests understanding of properties of buffers. The carbonic acid buffer contains both a weak acid (H₂CO₃) and its conjugate base (HCO₃⁻), though with excess weak acid present. When HCl is added, the H⁺ ions are primarily consumed by the bicarbonate ions (HCO₃⁻) through the reaction: HCO₃⁻ + H⁺ → H₂CO₃. This neutralization prevents most of the added H⁺ from remaining free in solution, resulting in only a slight pH decrease rather than a sharp drop. Option C incorrectly suggests that buffers with excess weak acid cannot react with added H⁺—this misunderstands that it's the conjugate base component that neutralizes added acid. The strategy for buffer problems is to identify which component reacts with the added species: conjugate bases neutralize added acids.
A buffer is prepared using the weak base $\mathrm{B}$ and its conjugate acid $\mathrm{BH^+}$, with $\mathrm{BH^+}$ present in excess. A small amount of $\mathrm{NaOH(aq)}$ is added. Which statement best describes what happens?
The pH decreases slightly because $\mathrm{OH^-}$ reacts with $\mathrm{B}$ to form $\mathrm{BH^+}$.
The pH increases only slightly because $\mathrm{OH^-}$ is consumed by $\mathrm{BH^+}$ to form $\mathrm{B}$ and $\mathrm{H_2O}$.
The pH decreases sharply because adding $\mathrm{OH^-}$ increases $[\mathrm{H_3O^+}]$ through water autoionization.
The pH increases sharply because the buffer contains mostly $\mathrm{BH^+}$ and cannot react with $\mathrm{OH^-}$.
The pH remains exactly constant because $\mathrm{BH^+}$ and $\mathrm{B}$ are both weak and therefore unreactive.
Explanation
This question tests understanding of properties of buffers. The buffer contains a weak base (B) and its conjugate acid (BH⁺), with excess BH⁺ present, allowing it to resist pH changes. When NaOH is added, the OH⁻ ions are consumed by the conjugate acid component (BH⁺) according to: BH⁺ + OH⁻ → B + H₂O. This neutralization reaction prevents most of the added hydroxide from remaining free in solution, resulting in only a slight pH increase rather than a sharp rise. Option A incorrectly suggests that OH⁻ reacts with B to form BH⁺—this is impossible as bases cannot react with other bases to form acids without a proton source. The strategy for buffer problems is to identify that conjugate acids neutralize added bases, while weak bases neutralize added acids.
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?
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.
The added $\mathrm{H^+}$ is primarily consumed by $\mathrm{HC_2H_3O_2}$ to form $\mathrm{C_2H_3O_2^-}$, so the pH remains exactly constant.
The added $\mathrm{H^+}$ reacts mainly with $\mathrm{Na^+}$ ions, preventing any change in pH.
The added $\mathrm{H^+}$ remains mostly as free $\mathrm{H^+}$ because buffers neutralize only added base, so the pH decreases sharply.
The added $\mathrm{H^+}$ causes $\mathrm{HC_2H_3O_2}$ to dissociate completely, making the solution behave like a strong acid.
Explanation
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 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?
The pH decreases sharply because $\mathrm{HA}$ is a weak acid and cannot participate in neutralization reactions.
The pH decreases only slightly because most added $\mathrm{H^+}$ is consumed by $\mathrm{A^-}$ to form $\mathrm{HA}$.
The pH remains exactly constant because the buffer completely prevents any change in $[\mathrm{H_3O^+}]$.
The pH increases because added $\mathrm{H^+}$ converts $\mathrm{HA}$ into $\mathrm{A^-}$.
The pH decreases sharply because the presence of excess $\mathrm{A^-}$ increases the concentration of free $\mathrm{H^+}$.
Explanation
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.
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?
The claim is incorrect; buffers work only when strong acids are added, not when strong bases are added.
The claim is incorrect; a buffer contains no species that can react with added $\mathrm{OH^-}$.
The claim is incorrect; added $\mathrm{OH^-}$ reacts with $\mathrm{HA}$ to form $\mathrm{A^-}$, so pH changes slightly rather than not at all.
The claim is correct; added $\mathrm{OH^-}$ reacts primarily with $\mathrm{A^-}$ to form $\mathrm{HA}$, leaving pH unchanged.
The claim is correct; buffers maintain a perfectly constant pH regardless of what is added.
Explanation
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 by mixing acetic acid, $\text{HC}_2\text{H}_3\text{O}_2(aq)$, and sodium acetate, $\text{NaC}_2\text{H}_3\text{O}_2(aq)$, so that the solution contains comparable amounts of $\text{HC}_2\text{H}_3\text{O}_2$ and $\text{C}_2\text{H}_3\text{O}_2^-$. A small amount of $\text{HCl}(aq)$ is added. Which statement best describes what happens in the solution?
The $\text{HCl}$ converts the buffer into a strong acid solution, causing a large decrease in pH.
The $\text{Na}^+$ ions react with the added $\text{H}^+$ to form $\text{NaH}(aq)$, preventing any pH change.
The added $\text{H}^+$ reacts primarily with water to form $\text{H}_3\text{O}^+$, so the pH decreases by the same amount as in pure water.
The added $\text{H}^+$ is consumed primarily by $\text{C}_2\text{H}_3\text{O}_2^-$ to form $\text{HC}_2\text{H}_3\text{O}_2$, so the pH decreases only slightly.
The added $\text{H}^+$ is consumed primarily by $\text{HC}_2\text{H}_3\text{O}_2$ to form $\text{H}_2\text{C}_2\text{H}_3\text{O}_2^+$, so the pH remains exactly constant.
Explanation
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 keep the pH exactly constant, as in choice B, but actually, small pH changes do occur, though they are minimized. 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 $\text{H}_2\text{CO}_3(aq)$ and $\text{HCO}_3^-(aq)$. The solution is then diluted by adding a large amount of pure water, with no acid or base added. Which statement best describes the effect on the buffer’s pH?
The pH decreases sharply because dilution always makes solutions more acidic.
The pH increases sharply because dilution always makes solutions more basic.
The pH becomes 7 because adding water forces neutrality.
The pH remains approximately the same because dilution lowers both buffer component concentrations proportionally, leaving their relative amounts unchanged.
The pH changes unpredictably because buffers only work at high concentration.
Explanation
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 dilution changes buffer pH like it does for strong acids, but as in choices B and C, buffer pH is stable due to the maintained ratio. To solve buffer problems, recall that pH depends on the ratio of components, so proportional changes like dilution do not significantly alter pH.