AP Chemistry Quiz: Properties Of Buffers

What this quiz covers

This quiz focuses on Properties Of Buffers, giving you a quick way to practice the rules, question types, and explanations that matter most for AP Chemistry.

How to use this quiz

Try each quiz question before looking at the correct answer. Use the explanations to review missed ideas, then come back to similar questions until the pattern feels familiar.

Question 1

1. The added base is neutralized mainly by $\mathrm{CHO_2^-}$ forming $\mathrm{HCHO_2}$, so the pH decreases slightly.
2. The pH stays exactly the same because the buffer converts all added base into water with no change in buffer composition.
3. The added base is neutralized mainly by $\mathrm{HCHO_2}$ forming $\mathrm{CHO_2^-}$ and water, so $\mathrm{[OH^-]}$ increases only slightly. (correct answer)
4. The pH changes little because $\mathrm{HCHO_2}$ is a strong acid that completely neutralizes any added base.
5. The pH changes little because $\mathrm{Na^+}$ consumes $\mathrm{OH^-}$ to form $\mathrm{NaOH(aq)}$, removing base from solution.

Explanation: This question tests understanding of buffer behavior when strong base is added to a formic acid/formate buffer. When strong base is added, the OH⁻ ions react with formic acid (HCHO₂) to form formate ions (CHO₂⁻) and water: HCHO₂ + OH⁻ → CHO₂⁻ + H₂O. This reaction neutralizes most of the added base by converting it to water and the conjugate base, so the [OH⁻] increases only slightly, minimizing pH change. Option A incorrectly suggests the base reacts with CHO₂⁻ to form HCHO₂, which would require protonating an anion with hydroxide ions, violating basic chemistry principles. The strategy is to recognize that added base reacts with the acidic component (the weak acid) to form its conjugate base and water.

Question 2

1. The pH does not change because $\mathrm{HF}$ is a strong acid and overwhelms the effect of adding $\mathrm{HBr}$.
2. The pH changes little because $\mathrm{NaF}$ is a strong base that neutralizes all added $\mathrm{HBr}$ completely.
3. The pH stays exactly constant because $\mathrm{Br^-}$ reacts with water to remove $\mathrm{H^+}$ from solution.
4. The added $\mathrm{H^+}$ is consumed primarily by $\mathrm{F^-}$ to form $\mathrm{HF}$, so the pH decreases only slightly. (correct answer)
5. The added $\mathrm{H^+}$ is consumed primarily by $\mathrm{HF}$ to form $\mathrm{F^-}$, which prevents a pH decrease.

Explanation: This question tests understanding of buffer action when strong acid is added to a hydrofluoric acid/fluoride buffer. When HBr is added, the H⁺ ions react with fluoride ions (F⁻) to form more hydrofluoric acid (HF): H⁺ + F⁻ → HF. This reaction consumes most of the added H⁺, converting it to the weak acid HF, which only partially dissociates, resulting in only a slight pH decrease. Option C incorrectly suggests H⁺ reacts with HF to form F⁻, which would require deprotonating an acid with additional acid, violating basic acid-base chemistry. The key strategy is to identify that added acid reacts with the basic component of the buffer (the conjugate base F⁻) to form the weak acid.

Question 3

1. The added acid reacts mainly with $\mathrm{H_2CO_3}$ to produce $\mathrm{CO_3^{2-}}$, so the buffer resists pH change.
2. The added acid reacts mainly with $\mathrm{NO_3^-}$ to form $\mathrm{HNO_3}$, which keeps the pH nearly constant.
3. The added acid reacts mainly with $\mathrm{HCO_3^-}$ to form $\mathrm{H_2CO_3}$, consuming most added $\mathrm{H^+}$ so pH changes only slightly. (correct answer)
4. The pH does not change because $\mathrm{H_2CO_3}$ and $\mathrm{HCO_3^-}$ completely neutralize any strong acid added.
5. The added acid reacts mainly with water to produce $\mathrm{OH^-}$, which offsets the added $\mathrm{H^+}$ in the buffer.

Explanation: This question tests understanding of buffer behavior in the carbonic acid/bicarbonate system when strong acid is added. When HNO₃ is added to the buffer, the H⁺ ions react primarily with bicarbonate ions (HCO₃⁻) to form carbonic acid (H₂CO₃): H⁺ + HCO₃⁻ → H₂CO₃. This reaction consumes most of the added H⁺, converting it to the weak acid H₂CO₃, which only partially dissociates, thus the pH changes only slightly. Option A incorrectly suggests H⁺ reacts with H₂CO₃ to form CO₃²⁻, which would require removing protons from an acid rather than adding them. The key strategy is to identify that added acid reacts with the basic component of the buffer (the conjugate base HCO₃⁻) to form the weak acid.

Question 4

1. The added $\mathrm{H^+}$ is consumed primarily by $\mathrm{C_2H_3O_2^-}$ to form $\mathrm{HC_2H_3O_2}$, removing most of the added acid so the pH changes only slightly. (correct answer)
2. The buffer prevents any pH change because weak acids and their salts completely neutralize strong acids as long as any solute remains.
3. The added $\mathrm{H^+}$ is consumed primarily by $\mathrm{HC_2H_3O_2}$ to form $\mathrm{C_2H_3O_2^-}$, so the pH stays constant.
4. The pH changes little because $\mathrm{NaC_2H_3O_2}$ makes the solution a strong base, which neutralizes $\mathrm{HCl}$ completely and permanently.
5. The pH changes little because $\mathrm{Cl^-}$ from $\mathrm{HCl}$ reacts with water to form $\mathrm{OH^-}$, offsetting the added $\mathrm{H^+}$.

Explanation: This question tests understanding of how buffers resist pH changes when strong acid is added. When HCl is added to an acetic acid/acetate buffer, the H⁺ ions from the strong acid are consumed by the acetate ions (C₂H₃O₂⁻) to form more acetic acid (HC₂H₃O₂), according to the reaction: H⁺ + C₂H₃O₂⁻ → HC₂H₃O₂. This reaction removes most of the added H⁺ from solution, converting it to the weak acid form, which only partially dissociates, thus minimizing the pH change. Option B incorrectly suggests buffers completely prevent any pH change, which violates the principle that buffers only minimize, not eliminate, pH changes. The key strategy is to identify which buffer component (the conjugate base) reacts with added acid to form the weak acid.

Question 5

1. The added $\mathrm{H^+}$ reacts mainly with $\mathrm{HClO}$ to form $\mathrm{ClO^-}$, so the pH does not decrease.
2. The added $\mathrm{H^+}$ reacts mainly with $\mathrm{ClO^-}$ to form $\mathrm{HClO}$, so most added acid is converted to a weak acid and pH changes slightly. (correct answer)
3. The pH does not change because $\mathrm{NaClO}$ makes the solution strongly basic, preventing any effect from added acid.
4. The pH remains constant because buffers eliminate all added $\mathrm{H^+}$ without changing the amounts of buffer components.
5. The pH changes little because $\mathrm{Cl^-}$ formed from the acid reacts with water to produce $\mathrm{OH^-}$ that neutralizes $\mathrm{H^+}$.

Explanation: This question tests understanding of buffer action when strong acid is added to a hypochlorous acid/hypochlorite system. When strong acid is added, the H⁺ ions react with hypochlorite ions (ClO⁻) to form hypochlorous acid (HClO): H⁺ + ClO⁻ → HClO. This reaction consumes most of the added H⁺, converting it to the weak acid HClO, which only partially dissociates, thus the pH changes only slightly. Option B incorrectly suggests H⁺ reacts with HClO to form ClO⁻, which would require deprotonating an acid using additional acid, a chemical impossibility. The key strategy is to identify that added acid reacts with the basic component of the buffer (the conjugate base ClO⁻) to form the weak acid.

Question 6

1. The pH changes little because the solution is effectively a strong acid due to $\mathrm{HNO_2}$, so added base is fully neutralized.
2. The pH stays exactly constant because buffers prevent any change in $\mathrm{[H^+]}$ regardless of added base amount.
3. The pH is unchanged because $\mathrm{K^+}$ reacts with $\mathrm{HNO_2}$ to form a neutral salt that fixes the pH.
4. The added $\mathrm{OH^-}$ is consumed primarily by $\mathrm{NO_2^-}$ to form $\mathrm{HNO_2}$, limiting the pH increase.
5. The added $\mathrm{OH^-}$ is consumed primarily by $\mathrm{HNO_2}$ to form $\mathrm{NO_2^-}$ and $\mathrm{H_2O}$, limiting the pH increase. (correct answer)

Explanation: This question tests understanding of how buffers respond to added strong base in a nitrous acid/nitrite system. When KOH is added, the OH⁻ ions react with the weak acid HNO₂ to form nitrite ions (NO₂⁻) and water: HNO₂ + OH⁻ → NO₂⁻ + H₂O. This reaction consumes the added hydroxide ions, converting them to water and the conjugate base NO₂⁻, which limits the pH increase to a small amount. Option B incorrectly suggests OH⁻ reacts with NO₂⁻ to form HNO₂, which would require protonating a base with hydroxide, a chemically impossible reaction. The strategy is to recognize that added base always reacts with the acidic component (the weak acid) in the buffer system.

Question 7

1. The pH stays exactly the same because buffers neutralize all added base regardless of the amount added.
2. The added $\mathrm{OH^-}$ is consumed primarily by $\mathrm{NH_4^+}$ to form $\mathrm{NH_3}$ and $\mathrm{H_2O}$, so the pH increases only slightly. (correct answer)
3. The added $\mathrm{OH^-}$ is consumed primarily by $\mathrm{NH_3}$ to form $\mathrm{NH_4^+}$, so the pH decreases slightly.
4. The pH changes little because $\mathrm{NH_4Cl}$ is a strong acid that neutralizes $\mathrm{OH^-}$ completely, leaving no excess base.
5. The pH changes little because $\mathrm{Na^+}$ reacts with water to form $\mathrm{H^+}$, counteracting the added $\mathrm{OH^-}$.

Explanation: This question tests understanding of buffer action when strong base is added to an ammonia/ammonium buffer system. When NaOH is added, the OH⁻ ions react with the ammonium ions (NH₄⁺) to form ammonia (NH₃) and water: NH₄⁺ + OH⁻ → NH₃ + H₂O. This reaction consumes most of the added hydroxide ions, converting them to water and the weak base NH₃, which only partially accepts protons, resulting in only a slight pH increase. Option C incorrectly reverses the reaction, suggesting OH⁻ reacts with NH₃ to form NH₄⁺, which would actually decrease pH rather than increase it. The strategy is to recognize that added base reacts with the acidic component of the buffer (the conjugate acid) to form the weak base.

Question 8

1. The added $\mathrm{H^+}$ is consumed primarily by $\mathrm{CH_3NH_2}$ to form $\mathrm{CH_3NH_3^+}$, reducing the increase in $\mathrm{[H^+]}$. (correct answer)
2. The added $\mathrm{H^+}$ is consumed primarily by $\mathrm{CH_3NH_3^+}$ to form $\mathrm{CH_3NH_2}$, reducing the increase in $\mathrm{[H^+]}$.
3. The pH stays exactly constant because buffers neutralize all added strong acid without any limit.
4. The pH changes little because the buffer is essentially a strong base solution, so $\mathrm{HCl}$ is completely removed.
5. The pH changes little because $\mathrm{Cl^-}$ reacts with $\mathrm{H^+}$ to form $\mathrm{HCl}$, keeping $\mathrm{[H^+]}$ low.

Explanation: This question tests understanding of buffer behavior in a methylamine/methylammonium system when strong acid is added. When HCl is added, the H⁺ ions react with the weak base methylamine (CH₃NH₂) to form methylammonium ions (CH₃NH₃⁺): CH₃NH₂ + H⁺ → CH₃NH₃⁺. This reaction consumes most of the added H⁺, converting the weak base to its conjugate acid, which reduces the increase in [H⁺] and minimizes pH change. Option B incorrectly reverses the reaction, suggesting H⁺ reacts with CH₃NH₃⁺ to form CH₃NH₂, which would require removing a proton from a cation using acid. The strategy is to recognize that in basic buffers, added acid reacts with the weak base component to form its conjugate acid.

Question 9

1. The buffer prevents any pH change because weak bases do not react with strong acids, so added $\mathrm{HCl}$ stays as $\mathrm{H^+}$ without affecting pH.
2. The added $\mathrm{H^+}$ reacts primarily with $\mathrm{CH_3NH_2}$ to form $\mathrm{CH_3NH_3^+}$, so most added acid is converted to the conjugate acid and the pH decreases only slightly. (correct answer)
3. Because the buffer contains an amine, the solution acts like a strong base; therefore, added $\mathrm{HCl}$ is completely neutralized with no change in the acid–base equilibrium.
4. The added $\mathrm{H^+}$ reacts primarily with $\mathrm{CH_3NH_3^+}$ to form $\mathrm{CH_3NH_2}$, so the conjugate acid neutralizes strong acid and prevents a pH decrease.
5. The added $\mathrm{H^+}$ reacts primarily with $\mathrm{NO_3^-}$ to form $\mathrm{HNO_3}$, which removes $\mathrm{H^+}$ from solution and maintains the original pH.

Explanation: The skill examined is properties of buffers, focusing on how a basic methylamine-methylammonium buffer resists pH decreases from added strong acid. When HCl is added to the CH3NH2 and CH3NH3NO3 buffer, the H+ reacts primarily with the weak base CH3NH2 to form CH3NH3+. This reaction consumes most added H+, slightly changing the [CH3NH2]/[CH3NH3+] ratio but keeping the pH drop small. The buffer functions by transforming strong acid into the conjugate acid, minimizing excess H+. Choice B tempts by suggesting H+ reacts with CH3NH3+ to form CH3NH2, which misrepresents the roles and reflects confusion between acid and base buffer behaviors. Always identify the weak base in basic buffers, as it directly protonates to handle added H+.

Question 10

1. The added H$^+$ is consumed mainly by CH$_3$NH$_2$ to form CH$_3$NH$_3^+$; converting strong acid into the conjugate acid limits the pH decrease. (correct answer)
2. The added H$^+$ is consumed mainly by CH$_3$NH$_3^+$, which donates OH$^-$ to neutralize the acid and keep pH constant.
3. The added HCl is consumed mainly by Cl$^-$, forming Cl$_2$ and removing H$^+$ so the pH does not change.
4. The buffer prevents pH change because CH$_3$NH$_2$ is a strong base, so added HCl has no effect on pH.
5. The added H$^+$ is consumed mainly by water, and the buffer resists pH change because water is amphoteric.

Explanation: This question tests understanding of buffer action when acid is added to a methylamine/methylammonium system. When HCl is added to a CH₃NH₂/CH₃NH₃⁺ buffer, the methylamine (CH₃NH₂) acts as a base and reacts with H⁺ to form methylammonium ion: CH₃NH₂ + H⁺ → CH₃NH₃⁺. This reaction converts the strong acid into the conjugate acid of a weak base, which establishes an equilibrium rather than fully dissociating, thereby limiting the pH decrease. Choice B is incorrect because it claims CH₃NH₃⁺ donates OH⁻, which reveals the misconception that cations can be hydroxide donors—the ammonium ion is actually an acid. The strategy for amine buffers is to recognize that the free base (CH₃NH₂) will consume added H⁺.

Question 11

1. The added HCl is neutralized mainly by HC$_2$H$_3$O$_2$, which donates OH$^-$ to keep the pH constant, so the pH does not change.
2. The added HCl is consumed mainly by C$_2$H$_3$O$_2^-$, forming HC$_2$H$_3$O$_2$; converting strong acid to weak acid limits the pH decrease. (correct answer)
3. The buffer works because acetic acid is a strong acid and therefore overwhelms the added HCl, preventing any meaningful pH change.
4. The added HCl is consumed mainly by Na$^+$, forming NaCl and removing H$^+$ from solution, so the pH stays exactly the same.
5. The added HCl is consumed mainly by water to form H$_3$O$^+$, and the buffer resists pH change because water is present in large excess.

Explanation: This question tests understanding of how buffers resist pH changes when acid is added. When HCl (a strong acid) is added to an acetic acid/acetate buffer, the acetate ion (C₂H₃O₂⁻) acts as a base and reacts with the H⁺ from HCl to form acetic acid (HC₂H₃O₂): C₂H₃O₂⁻ + H⁺ → HC₂H₃O₂. This reaction converts the strong acid (HCl) into a weak acid (acetic acid), which only partially dissociates, thereby limiting the increase in H⁺ concentration and minimizing the pH decrease. Choice A is incorrect because it claims acetic acid donates OH⁻, which reveals the misconception that weak acids produce hydroxide ions rather than accepting protons through their conjugate base. The key strategy is to identify which buffer component (the conjugate base) will react with added acid to form the weak acid, thereby consuming most of the added H⁺.

Question 12

1. The added NaOH is neutralized mainly by Cl$^-$ to form HCl, which keeps the pH constant by replacing the base with acid.
2. The added OH$^-$ is consumed mainly by water, and the buffer resists pH change because dilution by water dominates the chemistry.
3. The added OH$^-$ is consumed mainly by NH$_3$, which acts as a strong acid and neutralizes all base so the pH does not change.
4. The buffer prevents any pH change because both NH$_3$ and NH$_4^+$ completely dissociate, keeping [H$^+$] fixed.
5. The added OH$^-$ is consumed mainly by NH$_4^+$ to form NH$_3$ and H$_2$O; removing most of the added base limits the pH increase. (correct answer)

Explanation: This question tests understanding of buffer action when base is added to an ammonia/ammonium buffer system. When NaOH (providing OH⁻) is added to an NH₃/NH₄⁺ buffer, the ammonium ion (NH₄⁺) acts as an acid and reacts with OH⁻ to form ammonia and water: NH₄⁺ + OH⁻ → NH₃ + H₂O. This reaction consumes the added hydroxide ions by converting them to water, preventing a large increase in OH⁻ concentration and thus limiting the pH increase. Choice B is incorrect because it claims NH₃ acts as a strong acid, which reveals the misconception that ammonia (a weak base) can behave as a strong acid. The strategy for buffer problems with added base is to identify the acidic component of the buffer (here NH₄⁺) that will neutralize the added OH⁻.

Question 13

1. The added $\mathrm{H^+}$ reacts mainly with $\mathrm{ClO^-}$ to form $\mathrm{HClO}$, converting strong acid into weak acid so the pH changes only slightly. (correct answer)
2. The buffer works because $\mathrm{NaClO}$ is a strong base that completely neutralizes $\mathrm{HCl}$, so pH does not change.
3. The buffer prevents any pH change because $\mathrm{HClO}$ is a strong acid that overwhelms added $\mathrm{HCl}$ and fixes the pH.
4. The added $\mathrm{H^+}$ is neutralized mainly by $\mathrm{Na^+}$ ions, which react directly with $\mathrm{H^+}$ and keep pH constant.
5. The added $\mathrm{H^+}$ reacts mainly with $\mathrm{HClO}$ to form $\mathrm{ClO^-}$, so the buffer neutralizes acid by generating more base.

Explanation: This question tests understanding of how the HClO/ClO⁻ buffer responds to added acid. When HCl is added, the H⁺ ions react with the conjugate base ClO⁻ to form HClO through the reaction: ClO⁻ + H⁺ → HClO. This converts the strong acid HCl into the weak acid HClO, which only partially dissociates, thus limiting the pH change and maintaining the buffer's effectiveness. Choice C incorrectly reverses the reaction, suggesting that HClO consumes H⁺ to form ClO⁻, which would require the weak acid to act as a base and generate more base when acid is added. Remember that in any acid-base buffer pair, added H⁺ always reacts with the conjugate base to form more of the weak acid.

Question 14

1. The added $\mathrm{OH^-}$ is consumed mainly by $\mathrm{HC_3H_5O_3}$ to form $\mathrm{C_3H_5O_3^-}$ and water, so the conjugate acid/base ratio changes only slightly and the pH changes only slightly. (correct answer)
2. The added $\mathrm{OH^-}$ is consumed mainly by $\mathrm{C_3H_5O_3^-}$ to form $\mathrm{HC_3H_5O_3}$, so the conjugate base neutralizes base by becoming protonated.
3. The buffer prevents any pH change because the weak acid and its salt neutralize all added $\mathrm{OH^-}$ completely, leaving pH unchanged.
4. The buffer works because $\mathrm{NaC_3H_5O_3}$ acts as a strong base and directly neutralizes $\mathrm{NaOH}$, keeping the pH constant.
5. The added $\mathrm{OH^-}$ does not react in a buffer because weak acids do not donate protons readily, so the pH is unaffected by added base.

Explanation: This question tests understanding of buffer properties when base is added to a lactic acid/lactate buffer. When NaOH (strong base) is added to an HC₃H₅O₃/C₃H₅O₃⁻ buffer, the OH⁻ ions react with the weak acid (HC₃H₅O₃) to form the conjugate base (C₃H₅O₃⁻) and water: HC₃H₅O₃ + OH⁻ → C₃H₅O₃⁻ + H₂O. This neutralization reaction consumes the added OH⁻, preventing a large pH increase because the strong base is converted to the weakly basic lactate ion. The ratio of [C₃H₅O₃⁻]/[HC₃H₅O₃] increases slightly, causing only a small pH change according to the Henderson-Hasselbalch equation. Choice B incorrectly claims that OH⁻ reacts with C₃H₅O₃⁻ to form HC₃H₅O₃, which would require the conjugate base to become protonated by a base, violating fundamental acid-base principles since bases remove protons, not donate them. When strong base is added to any buffer, the weak acid component always neutralizes the OH⁻ through proton donation.

Question 15

1. The added $\mathrm{H^+}$ is consumed mainly by $\mathrm{HCOO^-}$ to form $\mathrm{HCOOH}$, so the buffer components remain in comparable amounts and the pH changes only slightly. (correct answer)
2. The added $\mathrm{H^+}$ is consumed mainly by $\mathrm{HCOOH}$ to form $\mathrm{HCOO^-}$, so adding acid increases the conjugate base and stabilizes the pH.
3. The buffer prevents any pH change because $\mathrm{HCOOH}$ is strong enough to keep $\mathrm{[H^+]}$ constant even when more acid is added.
4. The buffer works because $\mathrm{HCOONa}$ provides $\mathrm{OH^-}$ that neutralizes added acid completely, so the pH does not change.
5. The added $\mathrm{H^+}$ stays mostly unreacted because $\mathrm{HCOO^-}$ is too weak a base to accept protons, so the pH is unaffected.

Explanation: This question tests understanding of buffer action when acid is added to a formic acid/formate buffer system. When strong acid is added to an HCOOH/HCOO⁻ buffer, the H⁺ ions are consumed by the conjugate base (HCOO⁻) to form the weak acid (HCOOH): HCOO⁻ + H⁺ → HCOOH. This reaction prevents a large pH decrease by removing most of the added H⁺ from solution and converting it to the weak acid HCOOH, which only partially dissociates. The ratio of [HCOO⁻]/[HCOOH] decreases slightly, resulting in only a small pH change. Choice B incorrectly claims that H⁺ reacts with HCOOH to form HCOO⁻, which would require the acid to accept a proton and lose one simultaneously, a chemical impossibility that violates acid-base principles. To identify which buffer component reacts with added acid, look for the conjugate base (the species with one less H⁺) as the proton acceptor.

Question 16

1. The buffer prevents any pH change because $\mathrm{CH_3NH_2}$ is a strong base that completely neutralizes any added $\mathrm{H^+}$.
2. The added $\mathrm{H^+}$ is consumed mainly by $\mathrm{CH_3NH_2}$ to form $\mathrm{CH_3NH_3^+}$, so the conjugate base/acid ratio changes only slightly and the pH changes only slightly. (correct answer)
3. The added $\mathrm{H^+}$ does not react significantly in a buffer because weak bases do not accept protons readily, so the pH is unchanged.
4. The added $\mathrm{H^+}$ is consumed mainly by $\mathrm{CH_3NH_3^+}$ to form $\mathrm{CH_3NH_2}$, so the conjugate acid neutralizes added acid by losing a proton.
5. The buffer works because $\mathrm{CH_3NH_3Cl}$ is a strong acid that reacts completely with $\mathrm{H^+}$, so the pH stays constant.

Explanation: This question tests understanding of buffer properties when acid is added to a methylamine/methylammonium buffer. When strong acid is added to a CH₃NH₂/CH₃NH₃⁺ buffer, the H⁺ ions are consumed by the weak base (CH₃NH₂) to form the conjugate acid (CH₃NH₃⁺): CH₃NH₂ + H⁺ → CH₃NH₃⁺. This reaction prevents a large pH decrease by removing the added H⁺ from solution through protonation of the weak base. The ratio of [CH₃NH₂]/[CH₃NH₃⁺] decreases slightly, causing only a small pH change. Choice B incorrectly states that H⁺ reacts with CH₃NH₃⁺ to form CH₃NH₂, which would require the conjugate acid to accept another proton while simultaneously losing one, violating basic acid-base chemistry principles. When strong acid is added to a basic buffer, the weak base component (the unprotonated species) always acts as the proton acceptor.

Question 17

1. The added base reacts mainly with $\mathrm{H_2PO_4^-}$ to form $\mathrm{HPO_4^{2-}}$ and water, so the pH rises only slightly. (correct answer)
2. The added base reacts mainly with $\mathrm{HPO_4^{2-}}$ to form $\mathrm{H_2PO_4^-}$ and water, so the pH rises only slightly.
3. The pH does not change because phosphate buffers keep $\mathrm{[H^+]}$ fixed regardless of how much base is added.
4. The pH changes little because $\mathrm{HPO_4^{2-}}$ is a strong base and neutralizes all added base by itself.
5. The pH changes little because the added base reacts with water to create $\mathrm{H^+}$, offsetting the added $\mathrm{OH^-}$.

Explanation: This question tests understanding of buffer action in the dihydrogen phosphate/hydrogen phosphate system when strong base is added. When strong base is added, the OH⁻ ions react with H₂PO₄⁻ (which can act as an acid) to form HPO₄²⁻ and water: H₂PO₄⁻ + OH⁻ → HPO₄²⁻ + H₂O. This reaction consumes the added hydroxide ions, converting them to water while shifting the buffer equilibrium, resulting in only a slight pH rise. Option B incorrectly suggests OH⁻ reacts with HPO₄²⁻ to form H₂PO₄⁻, which would require protonating a base using hydroxide ions, a chemically impossible process. The key strategy is to identify which buffer component can donate a proton (act as an acid) when base is added.

Question 18

1. The added $\mathrm{H^+}$ reacts mainly with $\mathrm{HCOO^-}$ to form $\mathrm{HCOOH}$, so the conjugate base removes most added acid and the pH decreases only slightly. (correct answer)
2. Because $\mathrm{HCl}$ is strong, it completely determines the pH after addition, and the $\mathrm{HCOOH/HCOO^-}$ pair does not react significantly.
3. The added $\mathrm{H^+}$ reacts mainly with $\mathrm{Na^+}$ to form $\mathrm{NaH}$, removing acid and keeping the formic acid equilibrium unchanged.
4. The buffer prevents any pH change because it contains equal amounts of $\mathrm{HCOOH}$ and $\mathrm{HCOO^-}$, so added $\mathrm{H^+}$ cannot alter the pH at all.
5. The added $\mathrm{H^+}$ reacts mainly with $\mathrm{HCOOH}$ to form $\mathrm{HCOO^-}$, so the weak acid neutralizes strong acid and prevents a pH decrease.

Explanation: This question evaluates knowledge of buffer properties, focusing on how a formic acid-formate buffer mitigates pH drops from small additions of strong acid. The H+ from added HCl reacts mainly with HCOO- to form HCOOH in this HCOOH and HCOONa buffer. This reaction neutralizes most H+, causing a small change in the [HCOO-]/[HCOOH] ratio and thus a slight pH decrease. The underlying principle is the buffer's ability to absorb acid without significant dissociation changes. Choice B is a tempting wrong answer, stating H+ reacts with HCOOH to form HCOO-, which inverts the roles and shows misunderstanding of acid addition in acidic buffers. For effective problem-solving, always specify that added acid protonates the conjugate base in acidic buffers.

Question 19

1. The added H$^+$ is consumed mainly by water, and the buffer effect is due to dilution rather than acidbase reactions.
2. The added H$^+$ is consumed mainly by ClO$^-$ to form HClO; converting strong acid to weak acid limits the pH decrease. (correct answer)
3. The buffer works because HClO is a strong acid, so adding HCl does not change the pH significantly.
4. The added H$^+$ is consumed mainly by HClO, which neutralizes strong acids completely so the pH does not change.
5. The added H$^+$ is consumed mainly by Na$^+$ to form NaH, which removes H$^+$ and keeps the pH constant.

Explanation: This question tests understanding of buffer response to added acid in a hypochlorous acid/hypochlorite system. When HCl is added to an HClO/ClO⁻ buffer, the hypochlorite ion (ClO⁻) acts as a base and reacts with H⁺ from HCl to form HClO: ClO⁻ + H⁺ → HClO. This reaction converts the strong acid (HCl) into the weak acid (HClO), which only partially dissociates, thereby limiting the increase in H⁺ concentration and minimizing pH decrease. Choice D is incorrect because it claims HClO is a strong acid, which is a misconception—HClO is actually a weak acid with a Ka around 3 × 10⁻⁸. The key strategy for acid addition to buffers is to identify the basic component (conjugate base) that will consume the added H⁺.

Question 20

1. The added H$^+$ is consumed mainly by CN$^-$ to form HCN; removing most added H$^+$ limits the pH decrease. (correct answer)
2. The added H$^+$ is consumed mainly by HCN, which neutralizes HCl completely so the pH does not change at all.
3. The added HCl is consumed mainly by Na$^+$, forming NaCl and eliminating H$^+$ from solution.
4. The buffer works because HCN fully dissociates, making the solution a strong acid that is unaffected by added HCl.
5. The added H$^+$ is consumed mainly by water, and the buffer resists pH change because water is present in large excess.

Explanation: This question tests understanding of buffer action when acid is added to a hydrocyanic acid/cyanide system. When HCl is added to an HCN/CN⁻ buffer, the cyanide ion (CN⁻) acts as a base and reacts with H⁺ from HCl to form HCN: CN⁻ + H⁺ → HCN. This reaction converts the strong acid into the extremely weak acid HCN (Ka ≈ 6.2 × 10⁻¹⁰), effectively removing most of the added H⁺ from solution and limiting pH decrease. Choice D is incorrect because it claims HCN fully dissociates, making it a strong acid—this is a serious misconception as HCN is one of the weakest acids commonly encountered. The key strategy is recognizing that the conjugate base (CN⁻) will consume added H⁺ to form the weak acid.