Properties of Buffers - AP Chemistry

A buffer is made by combining aqueous ammonia, $\mathrm{NH_3}$ , and ammonium chloride, $\mathrm{NH_4Cl}$ . A small amount of strong base, $\mathrm{NaOH(aq)}$ , is added. Which statement best describes how the buffer resists a large pH change and identifies the species that reacts with the added base?
Answer: The added $\mathrm{OH^-}$ is consumed mainly by $\mathrm{NH_4^+}$ to form $\mathrm{NH_3}$ and water, so the conjugate acid/base ratio shifts only slightly and the pH changes only slightly. This question tests understanding of buffer properties when base is added to an ammonia/ammonium buffer system. When NaOH (strong base) is added to an NH₃/NH₄⁺ buffer, the OH⁻ ions react with the conjugate acid (NH₄⁺) to form the weak base (NH₃) and water: NH₄⁺ + OH⁻ → NH₃ + H₂O. This reaction consumes the added OH⁻, preventing a large pH increase because the strong base is converted to the weak base NH₃, which only partially accepts protons. The ratio of [NH₃]/[NH₄⁺] increases slightly, causing only a small pH change. Choice B incorrectly suggests that OH⁻ reacts with NH₃ to form NH₄⁺, which would require the base to donate OH⁻ or accept H⁺ from OH⁻, both of which are chemically impossible since OH⁻ is a base, not an acid. When strong base is added to a buffer, always identify the conjugate acid component (the protonated species) as the one that neutralizes the added OH⁻.
A buffer is made by mixing formic acid, $\mathrm{HCOOH}$ , and sodium formate, $\mathrm{HCOONa}$ . A small amount of $\mathrm{HCl(aq)}$ is added. Which statement best explains how the buffer resists pH change and identifies the reacting species?
Answer: 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. 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.
A buffer is prepared using carbonic acid, $\mathrm{H_2CO_3}$ , and sodium bicarbonate, $\mathrm{NaHCO_3}$ . A small amount of strong base is added. Which statement best explains how and why the buffer resists a large pH change and which component reacts with the added base?
Answer: The added $\mathrm{OH^-}$ is consumed mainly by $\mathrm{H_2CO_3}$ to form $\mathrm{HCO_3^-}$ and water, so the ratio of $\mathrm{H_2CO_3}$ to $\mathrm{HCO_3^-}$ changes only slightly and the pH changes only slightly. This question tests understanding of buffer properties when base is added to a carbonic acid/bicarbonate buffer. When strong base is added to an H₂CO₃/HCO₃⁻ buffer, the OH⁻ ions react with the weak acid (H₂CO₃) to form the conjugate base (HCO₃⁻) and water: H₂CO₃ + OH⁻ → HCO₃⁻ + H₂O. This neutralization consumes the added OH⁻, preventing a large pH increase because the strong base is converted to the weakly basic HCO₃⁻. The ratio of [HCO₃⁻]/[H₂CO₃] increases slightly, causing only a small pH change. Choice B incorrectly suggests that OH⁻ reacts with HCO₃⁻ to form H₂CO₃, which would require the conjugate base to gain a proton from hydroxide ion, an impossible reaction since OH⁻ is a base that accepts protons, not a proton donor. When analyzing buffer responses to added base, the weak acid component (the more protonated species) always acts as the proton donor to neutralize OH⁻.
A buffer is made by combining $\mathrm{HCN(aq)}$ and $\mathrm{CN^- (aq)}$ (from $\mathrm{NaCN}$ ). A small amount of $\mathrm{HCl(aq)}$ is added. Which statement best describes how the buffer resists pH change and which species reacts with the added acid?
Answer: The added $\mathrm{H^+}$ reacts mainly with $\mathrm{CN^-}$ to form $\mathrm{HCN}$ , so most of the strong acid is converted into weak acid and the pH changes only slightly. This question tests understanding of how the HCN/CN⁻ buffer responds to added acid. When HCl is added, the H⁺ ions react with the conjugate base CN⁻ to form HCN through the reaction: CN⁻ + H⁺ → HCN. This converts most of the strong acid into the weak acid HCN, which has limited dissociation, resulting in only a small pH change. Choice B incorrectly reverses the reaction, suggesting that HCN consumes H⁺ to form CN⁻, which would require the weak acid to act as a base and produce more base when acid is added. To identify the correct buffer component that reacts with added acid, always look for the conjugate base (the deprotonated form).
A buffer is made by combining aqueous ammonia, $\mathrm{NH_3}$ , and ammonium chloride, $\mathrm{NH_4Cl}$ . A small amount of $\mathrm{NaOH(aq)}$ is added. Which statement best describes how the buffer resists a large pH increase, including which component reacts with the added base?
Answer: The added $\mathrm{OH^-}$ reacts primarily with $\mathrm{NH_4^+}$ to form $\mathrm{NH_3}$ and $\mathrm{H_2O}$ , so the conjugate acid removes most of the added base and the pH changes only slightly. This question assesses knowledge of buffer properties, focusing on how a basic buffer resists pH increases upon addition of a small amount of strong base. In the ammonia and ammonium chloride buffer, the added NaOH provides OH- ions that react mainly with the conjugate acid, NH4+, to form NH3 and H2O. This neutralizes most of the added base, slightly shifting the [NH3]/[NH4+] ratio but keeping the pH change minimal according to the buffer equilibrium. The buffering action relies on the conjugate acid consuming OH- to prevent a large increase in pH. Choice D is a common distractor, wrongly claiming that the buffer prevents any pH change because NH3 is weak, based on the misconception that weak bases inherently block pH shifts without involving equilibrium reactions. A useful strategy is to remember that in basic buffers, added OH- is neutralized by the conjugate acid, forming more weak base and water.
A buffer contains hydrofluoric acid, $\mathrm{HF(aq)}$ , and sodium fluoride, $\mathrm{NaF(aq)}$ . A small amount of $\mathrm{HBr(aq)}$ is added. Which statement correctly describes what happens in the buffer?
Answer: The added $\mathrm{H^+}$ is consumed primarily by $\mathrm{F^-}$ to form $\mathrm{HF}$ , so the pH decreases only slightly. 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.
A buffer is made from $\mathrm{H_2CO_3(aq)}$ and $\mathrm{HCO_3^- (aq)}$ (from $\mathrm{NaHCO_3}$ ). A small amount of $\mathrm{HNO_3(aq)}$ is added. Which statement best explains the buffer action, including which component reacts with the added acid?
Answer: The added $\mathrm{H^+}$ is consumed mainly by $\mathrm{HCO_3^-}$ to form $\mathrm{H_2CO_3}$ , so the buffer components convert strong acid into weak acid and limit pH change. This question tests understanding of how the carbonic acid/bicarbonate buffer responds to added acid. When HNO₃ is added, the H⁺ ions react with the conjugate base HCO₃⁻ to form H₂CO₃ through the reaction: HCO₃⁻ + H⁺ → H₂CO₃. This converts the strong acid HNO₃ into the weak acid H₂CO₃, which only partially dissociates, thus limiting the pH change. Choice B incorrectly reverses the reaction, suggesting that H₂CO₃ consumes H⁺ to form HCO₃⁻, which would require the weak acid to act as a base. Remember that in a buffer system, added acid reacts with the conjugate base component to form more of the weak acid.
A buffer contains HClO(aq) and ClO $^-$ (aq) (from NaClO). A small amount of HCl(aq) is added. Which statement best describes the buffer mechanism and identifies what reacts with the added acid?
Answer: The added H $^+$ is consumed mainly by ClO $^-$ to form HClO; converting strong acid to weak acid limits the pH decrease. 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⁺.
A buffer is prepared by mixing aqueous $\mathrm{CH_3COOH}$ and $\mathrm{CH_3COO^-}$ (from $\mathrm{NaCH_3COO}$ ). A small amount of $\mathrm{HCl(aq)}$ is added to the buffer. Which statement best explains how and why the buffer resists a large pH change, including which component reacts with the added acid?
Answer: Most added $\mathrm{H^+}$ is consumed by $\mathrm{CH_3COO^-}$ to form $\mathrm{CH_3COOH}$ , so the ratio of conjugate base to acid changes only slightly and pH changes little. This question tests understanding of how buffers resist pH changes when acid is added. When HCl is added to an acetic acid/acetate buffer, the H⁺ ions react with the conjugate base CH₃COO⁻ to form more CH₃COOH, following the reaction: CH₃COO⁻ + H⁺ → CH₃COOH. This consumes most of the added H⁺, converting it from a strong acid into the weak acid CH₃COOH, which only slightly changes the ratio of conjugate base to acid and therefore causes minimal pH change. Choice C incorrectly reverses the reaction, suggesting that CH₃COOH consumes H⁺ to form CH₃COO⁻, which would require the weak acid to act as a base. When acid is added to a buffer, always identify the conjugate base component as the species that will neutralize the added H⁺.
A student makes a buffer by mixing aqueous formic acid, $\mathrm{HCOOH}$ , and sodium formate, $\mathrm{HCOO^-}$ (from $\mathrm{HCOONa}$ ). A small amount of strong acid is added. Which statement best describes the buffering mechanism and identifies which component reacts with the added acid?
Answer: 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. 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.
A buffer is prepared by mixing aqueous hydrofluoric acid, $\mathrm{HF}$ , with sodium fluoride, $\mathrm{NaF}$ . A small amount of strong acid is added. Which statement best explains the buffering action and identifies which component reacts with the added acid?
Answer: The added $\mathrm{H^+}$ is consumed mainly by $\mathrm{F^-}$ to form $\mathrm{HF}$ , so the relative amounts of $\mathrm{HF}$ and $\mathrm{F^-}$ change only slightly and the pH changes only slightly. This question tests understanding of buffer action when acid is added to a hydrofluoric acid/fluoride buffer. When strong acid is added to an HF/F⁻ buffer, the H⁺ ions are consumed by the conjugate base (F⁻) to form the weak acid (HF): F⁻ + H⁺ → HF. This reaction prevents a large pH decrease because it removes the added H⁺ from solution by converting it to HF, which as a weak acid only partially dissociates. The ratio of [F⁻]/[HF] decreases slightly, resulting in only a small pH change according to the Henderson-Hasselbalch equation. Choice B incorrectly states that H⁺ reacts with HF to form F⁻, which would require the acid to accept a proton and become deprotonated, contradicting the fundamental behavior of acids. When analyzing buffer responses to added acid, remember that the conjugate base component (the deprotonated species) always acts as the proton acceptor.
A buffer is prepared using H $_2$ CO $_3$ (aq) and NaHCO $_3$ (aq). A small amount of NaOH(aq) is added to the buffer. Which statement best explains how the buffer resists pH change and which component reacts with the added base?
Answer: The added OH $^-$ is consumed mainly by H $_2$ CO $_3$ to form HCO $_3^-$ and H $_2$ O; removing added base by forming the conjugate base limits the pH increase. This question tests understanding of buffer action in a carbonic acid/bicarbonate system when base is added. When NaOH (providing OH⁻) is added to an H₂CO₃/HCO₃⁻ buffer, the carbonic acid (H₂CO₃) acts as an acid and reacts with OH⁻ to form bicarbonate and water: H₂CO₃ + OH⁻ → HCO₃⁻ + H₂O. This reaction consumes the added base by converting it to water and the conjugate base (HCO₃⁻), preventing a large increase in OH⁻ concentration and limiting the pH increase. Choice A is incorrect because it suggests HCO₃⁻ reacts with OH⁻ to form CO₃²⁻, which would actually occur if excess base were added but doesn't represent the primary buffer action. The strategy is to identify the acidic component of the buffer pair (H₂CO₃) that will neutralize added base.
A buffer is prepared from HNO $_2$ (aq) and NaNO $_2$ (aq). A small amount of HCl(aq) is added. Which statement best describes how the buffer minimizes the pH change and which species reacts with the added acid?
Answer: The added H $^+$ is consumed mainly by NO $_2^-$ to form HNO $_2$ ; removing most added H $^+$ limits the pH decrease. This question tests understanding of buffer response to added acid in a nitrous acid/nitrite system. When HCl is added to an HNO₂/NO₂⁻ buffer, the nitrite ion (NO₂⁻) acts as a base and reacts with H⁺ from HCl to form HNO₂: NO₂⁻ + H⁺ → HNO₂. This reaction converts the strong acid into the weak acid HNO₂, which only partially dissociates, thereby removing most of the added H⁺ and limiting the pH decrease. Choice B is incorrect because it suggests HNO₂ reacts with H⁺ to form H₂NO₂⁺, which represents a misconception about how weak acids behave—they don't typically protonate further under normal conditions. The key strategy is recognizing that the conjugate base component (NO₂⁻) consumes added acid.
A buffer is prepared by mixing aqueous $HC$ _2 $H$ _3 $O$ _2$ (acetic acid) and aqueous NaC $_2$ H $_3$ O $_2$ (sodium acetate). A small amount of HCl(aq) is added to the buffer. Which statement best explains how and why the buffer resists a large change in pH, including which component reacts with the added acid?
Answer: 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. 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⁺.
A buffer is made from nitrous acid, $\mathrm{HNO_2(aq)}$ , and sodium nitrite, $\mathrm{NaNO_2(aq)}$ . A small amount of strong base, $\mathrm{KOH(aq)}$ , is added. Which statement best explains the buffer behavior?
Answer: The added $\mathrm{OH^-}$ is consumed primarily by $\mathrm{HNO_2}$ to form $\mathrm{NO_2^-}$ and $\mathrm{H_2O}$ , limiting the pH increase. 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.
A buffer is prepared by mixing benzoic acid, $\mathrm{HC_7H_5O_2}$ , and sodium benzoate, $\mathrm{NaC_7H_5O_2}$ . A small amount of $\mathrm{NaOH(aq)}$ is added. Which statement best explains how the buffer resists pH change and identifies which component reacts with the added base?
Answer: The added $\mathrm{OH^-}$ reacts mainly with $\mathrm{HC_7H_5O_2}$ to form $\mathrm{C_7H_5O_2^-}$ and $\mathrm{H_2O}$ , so the weak acid neutralizes most added base and the pH increases only slightly. This question assesses buffer properties, particularly how a benzoic acid-benzoate buffer counters pH increases from added strong base. The OH- from added NaOH reacts mainly with the weak acid HC7H5O2 to form C7H5O2- and H2O in this buffer. This neutralizes the base, causing a small shift in the [C7H5O2-]/[HC7H5O2] ratio and only a slight pH rise. The principle is that the weak acid donates H+ to OH-, converting it to water and the conjugate base. Choice A is incorrect and tempting, claiming OH- reacts with C7H5O2- to form HC7H5O2, which confuses the reaction direction and misapplies the buffering mechanism for bases. To tackle buffer problems, note that added OH- is neutralized by the acidic component, regardless of the buffer's pKa.
A buffer contains $\mathrm{HClO(aq)}$ and $\mathrm{ClO^-(aq)}$ (from $\mathrm{NaClO}$ ). If a small amount of strong acid is added, which statement correctly describes the buffering process?
Answer: 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. 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.
A student prepares a buffer by combining $\mathrm{H_2CO_3(aq)}$ and $\mathrm{NaHCO_3(aq)}$ . The student then adds a small amount of $\mathrm{HNO_3(aq)}$ . Which statement correctly identifies the species that reacts most directly with the added acid and explains the small pH change?
Answer: 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. 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.
A buffer is prepared by mixing hydrofluoric acid, $\mathrm{HF}$ , and sodium fluoride, $\mathrm{NaF}$ . A small amount of $\mathrm{NaOH(aq)}$ is added. Which statement correctly describes how the buffer resists pH change and which component reacts with the added base?
Answer: The added $\mathrm{OH^-}$ reacts primarily with $\mathrm{HF}$ to form $\mathrm{F^-}$ and $\mathrm{H_2O}$ , so the weak acid neutralizes most of the base and the pH increases only slightly. This question evaluates understanding of buffer properties, emphasizing how an HF-F- buffer resists pH increases when a small amount of strong base is added. The added NaOH introduces OH- ions that react predominantly with the weak acid HF to form F- and H2O. This neutralization consumes most OH-, slightly altering the [F-]/[HF] ratio but resulting in only a small pH increase per the Henderson-Hasselbalch principle. The buffer works by converting the strong base into the conjugate base, limiting free OH- in solution. A tempting incorrect choice is B, which reverses the roles by saying F- reacts with OH- to form HF, revealing the misconception of confusing which buffer component handles added base. To solve similar problems, recall that added base reacts with the acidic component of the buffer to form its conjugate and water.
A buffer contains hypochlorous acid, $\mathrm{HOCl}$ , and sodium hypochlorite, $\mathrm{NaOCl}$ . A small amount of $\mathrm{HCl(aq)}$ is added. Which statement best describes how the buffer resists a large pH decrease and which component reacts with the added acid?
Answer: The added $\mathrm{H^+}$ reacts mainly with $\mathrm{OCl^-}$ to form $\mathrm{HOCl}$ , so most added acid is converted to the weak acid form and the pH changes only slightly. The skill being tested is understanding buffer properties, specifically how a hypochlorous acid-hypochlorite buffer resists pH decreases when strong acid is added. In the HOCl and NaOCl buffer, added HCl's H+ reacts primarily with the conjugate base OCl- to form HOCl. This consumes the H+, slightly adjusting the [OCl-]/[HOCl] ratio but resulting in only a minor pH change. The buffer's resistance comes from converting strong acid to weak acid via equilibrium shift. Choice A distracts by reversing it to H+ reacting with HOCl to form OCl-, based on the misconception of which component absorbs added acid. A general strategy is to recall that conjugate bases in buffers protonate to handle added H+, preserving pH stability.

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