Acid-Base Reactions and Buffers
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AP Chemistry › Acid-Base Reactions and Buffers
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 remains exactly constant because buffers completely eliminate added $\mathrm{H^+}$.
The pH increases because adding $\mathrm{HCl}$ shifts $\mathrm{H_2CO_3}$ to produce more $\mathrm{HCO_3^-}$.
The added $\mathrm{H^+}$ is primarily consumed by $\mathrm{HCO_3^-}$ to form $\mathrm{H_2CO_3}$, so the pH decreases slightly.
The added $\mathrm{H^+}$ reacts mainly with $\mathrm{Na^+}$, so the pH changes very little.
The pH decreases sharply because the buffer contains excess weak acid and cannot react with 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 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 $\mathrm{H^+}$ reacts mainly with $\mathrm{Na^+}$ ions, preventing any change in pH.
The added $\mathrm{H^+}$ causes $\mathrm{HC_2H_3O_2}$ to dissociate completely, making the solution behave like a strong acid.
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 $\mathrm{H^+}$ is primarily consumed by $\mathrm{C_2H_3O_2^-}$ to form $\mathrm{HC_2H_3O_2}$, so the pH decreases only slightly.
The added $\mathrm{H^+}$ remains mostly as free $\mathrm{H^+}$ because buffers neutralize only added base, so the pH decreases sharply.
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 increases because added $\mathrm{H^+}$ converts $\mathrm{HA}$ into $\mathrm{A^-}$.
The pH decreases only slightly because most added $\mathrm{H^+}$ is consumed by $\mathrm{A^-}$ to form $\mathrm{HA}$.
The pH decreases sharply because the presence of excess $\mathrm{A^-}$ increases the concentration of free $\mathrm{H^+}$.
The pH remains exactly constant because the buffer completely prevents any change in $[\mathrm{H_3O^+}]$.
The pH decreases sharply because $\mathrm{HA}$ is a weak acid and cannot participate in neutralization reactions.
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 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 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 $\text{HC}_2\text{H}_3\text{O}_2$ to form $\text{H}_2\text{C}_2\text{H}_3\text{O}_2^+$, so the pH remains exactly constant.
The $\text{HCl}$ converts the buffer into a strong acid solution, causing a large decrease in pH.
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 $\text{Na}^+$ ions react with the added $\text{H}^+$ to form $\text{NaH}(aq)$, preventing any pH change.
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 becomes 7 because adding water forces neutrality.
The pH increases sharply because dilution always makes solutions more basic.
The pH decreases sharply because dilution always makes solutions more acidic.
The pH changes unpredictably because buffers only work at high concentration.
The pH remains approximately the same because dilution lowers both buffer component concentrations proportionally, leaving their relative amounts unchanged.
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.
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?
$\mathrm{HC_2H_3O_2}$ reacts with added $\mathrm{H^+}$ to form $\mathrm{H_2C_2H_3O_2^+}$, preventing any pH change.
The added $\mathrm{HCl}$ fully dissociates, so the pH decreases by the same amount as it would in pure water.
$\mathrm{C_2H_3O_2^-}$ reacts with added $\mathrm{H^+}$ to form $\mathrm{HC_2H_3O_2}$, removing most of the added $\mathrm{H^+}$.
$\mathrm{Na^+}$ reacts with added $\mathrm{H^+}$ to form $\mathrm{NaH(aq)}$, which neutralizes the acid.
The added $\mathrm{H^+}$ is consumed mainly by $\mathrm{OH^-}$ already present in the buffer, so the pH stays exactly constant.
Explanation
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 is prepared by mixing $\mathrm{H_2PO_4^-}$ and $\mathrm{HPO_4^{2-}}$ in water so that both species are present. A small amount of $\mathrm{HCl(aq)}$ is added. Which statement best describes the primary reaction that helps resist the pH change?
$\mathrm{Cl^-}$ reacts with water to form $\mathrm{HClO}$, consuming $\mathrm{H^+}$.
Water reacts with the added $\mathrm{H^+}$ to form $\mathrm{OH^-}$, which neutralizes the acid.
$\mathrm{H_2PO_4^-}$ reacts with the added $\mathrm{H^+}$ to form $\mathrm{HPO_4^{2-}}$, consuming $\mathrm{H^+}$.
The added $\mathrm{H^+}$ is consumed primarily by $\mathrm{Na^+}$ impurities in the solution, forming $\mathrm{NaH(aq)}$.
$\mathrm{HPO_4^{2-}}$ reacts with the added $\mathrm{H^+}$ to form $\mathrm{H_2PO_4^-}$, consuming $\mathrm{H^+}$.
Explanation
This question tests understanding of properties of buffers. The buffer contains the dihydrogen phosphate ion (H₂PO₄⁻) and hydrogen phosphate ion (HPO₄²⁻), which form a conjugate acid-base pair that resists pH changes. When HCl is added, the H⁺ ions react with the more basic species, HPO₄²⁻, to form H₂PO₄⁻ through the reaction: HPO₄²⁻ + H⁺ → H₂PO₄⁻. This reaction consumes the added H⁺ ions, preventing them from significantly lowering the pH of the solution. Choice A is incorrect because it shows H₂PO₄⁻ reacting with H⁺ to form HPO₄²⁻, which would require removing a proton from an already protonated species—this is the opposite of what happens when acid is added. The key strategy is to identify the more basic component in the buffer (the one with fewer protons), as this will be the species that reacts with added acid.
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?
The added $\mathrm{H^+}$ is primarily consumed by $\mathrm{CH_3NH_3^+}$ to form $\mathrm{CH_3NH_4^{2+}}$.
The added $\mathrm{H^+}$ is primarily consumed by $\mathrm{Cl^-}$ to form $\mathrm{HCl}$, which is weak.
The pH decreases more than in pure water because the buffer contains additional dissolved ions.
The pH does not change because $\mathrm{CH_3NH_2}$ is a strong base that neutralizes all added acid.
The added $\mathrm{H^+}$ is primarily consumed by $\mathrm{CH_3NH_2}$ to form $\mathrm{CH_3NH_3^+}$, reducing the amount of free $\mathrm{H^+}$.
Explanation
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.
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?
Solution 1 shows a smaller pH decrease because weak acids resist pH change better than buffers.
Solution 2 shows a smaller pH decrease because $\text{HCOO}^-$ consumes much of the added $\text{H}^+$.
Both solutions show the same pH decrease because the same amount of $\text{HCl}$ is added.
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.
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 $\mathrm{OH^-}$ reacts mainly with $\mathrm{Na^+}$, preventing a pH change.
The pH remains exactly constant because the buffer converts all added $\mathrm{OH^-}$ into neutral salt with no equilibrium shift.
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 $\mathrm{OH^-}$ is primarily neutralized by $\mathrm{F^-}$ to form $\mathrm{HF}$, so the pH decreases only slightly.
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.