Properties of Buffers
Help Questions
AP Chemistry › Properties of Buffers
A buffer is prepared by mixing aqueous formic acid, $\mathrm{HCHO_2(aq)}$, and sodium formate, $\mathrm{NaCHO_2(aq)}$. A small amount of strong base is added. Which statement best explains the resistance to pH change?
The added base is neutralized mainly by $\mathrm{CHO_2^-}$ forming $\mathrm{HCHO_2}$, so the pH decreases slightly.
The pH changes little because $\mathrm{HCHO_2}$ is a strong acid that completely neutralizes any added base.
The added base is neutralized mainly by $\mathrm{HCHO_2}$ forming $\mathrm{CHO_2^-}$ and water, so $\mathrm{[OH^-]}$ increases only slightly.
The pH stays exactly the same because the buffer converts all added base into water with no change in buffer composition.
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.
A buffer is prepared using $\mathrm{H_2PO_4^-(aq)}$ and $\mathrm{HPO_4^{2-}(aq)}$. A small amount of strong base is added. Which statement best describes the reaction that helps the buffer resist a large pH change?
The pH does not change because phosphate buffers keep $\mathrm{[H^+]}$ fixed regardless of how much base is added.
The pH changes little because the added base reacts with water to create $\mathrm{H^+}$, offsetting the added $\mathrm{OH^-}$.
The added base reacts mainly with $\mathrm{H_2PO_4^-}$ to form $\mathrm{HPO_4^{2-}}$ and water, so the pH rises only slightly.
The pH changes little because $\mathrm{HPO_4^{2-}}$ is a strong base and neutralizes all added base by itself.
The added base reacts mainly with $\mathrm{HPO_4^{2-}}$ to form $\mathrm{H_2PO_4^-}$ and water, so the pH rises only slightly.
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.
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?
The added $\mathrm{H^+}$ is consumed primarily by $\mathrm{HF}$ to form $\mathrm{F^-}$, which prevents a pH decrease.
The pH does not change because $\mathrm{HF}$ is a strong acid and overwhelms the effect of adding $\mathrm{HBr}$.
The added $\mathrm{H^+}$ is consumed primarily by $\mathrm{F^-}$ to form $\mathrm{HF}$, so the pH decreases only slightly.
The pH stays exactly constant because $\mathrm{Br^-}$ reacts with water to remove $\mathrm{H^+}$ from solution.
The pH changes little because $\mathrm{NaF}$ is a strong base that neutralizes all added $\mathrm{HBr}$ completely.
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.
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?
The pH does not change because $\mathrm{H_2CO_3}$ and $\mathrm{HCO_3^-}$ completely neutralize any strong acid added.
The added acid reacts mainly with $\mathrm{NO_3^-}$ to form $\mathrm{HNO_3}$, which keeps the pH nearly constant.
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.
The added acid reacts mainly with $\mathrm{H_2CO_3}$ to produce $\mathrm{CO_3^{2-}}$, so the buffer resists pH change.
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.
A buffer is prepared by mixing aqueous acetic acid, $\mathrm{HC_2H_3O_2}$, and sodium acetate, $\mathrm{NaC_2H_3O_2}$. A small amount of strong acid, $\mathrm{HCl(aq)}$, is added to the buffer. Which statement best explains how and why the buffer resists a large pH change?
The pH changes little because $\mathrm{NaC_2H_3O_2}$ makes the solution a strong base, which neutralizes $\mathrm{HCl}$ completely and permanently.
The buffer prevents any pH change because weak acids and their salts completely neutralize strong acids as long as any solute remains.
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.
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.
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.
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?
The added $\mathrm{H^+}$ reacts mainly with $\mathrm{HClO}$ to form $\mathrm{ClO^-}$, so the pH does not decrease.
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.
The pH remains constant because buffers eliminate all added $\mathrm{H^+}$ without changing the amounts of buffer components.
The pH does not change because $\mathrm{NaClO}$ makes the solution strongly basic, preventing any effect from added acid.
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.
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?
The added $\mathrm{OH^-}$ is consumed primarily by $\mathrm{HNO_2}$ to form $\mathrm{NO_2^-}$ and $\mathrm{H_2O}$, limiting the pH increase.
The added $\mathrm{OH^-}$ is consumed primarily by $\mathrm{NO_2^-}$ to form $\mathrm{HNO_2}$, limiting the pH increase.
The pH changes little because the solution is effectively a strong acid due to $\mathrm{HNO_2}$, so added base is fully neutralized.
The pH stays exactly constant because buffers prevent any change in $\mathrm{[H^+]}$ regardless of added base amount.
The pH is unchanged because $\mathrm{K^+}$ reacts with $\mathrm{HNO_2}$ to form a neutral salt that fixes the pH.
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.
A buffer contains ammonia, $\mathrm{NH_3(aq)}$, and ammonium chloride, $\mathrm{NH_4Cl(aq)}$. A small amount of strong base, $\mathrm{NaOH(aq)}$, is added. Which statement best describes the buffer action?
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.
The pH stays exactly the same because buffers neutralize all added base regardless of the amount added.
The pH changes little because $\mathrm{Na^+}$ reacts with water to form $\mathrm{H^+}$, counteracting the added $\mathrm{OH^-}$.
The pH changes little because $\mathrm{NH_4Cl}$ is a strong acid that neutralizes $\mathrm{OH^-}$ completely, leaving no excess base.
The added $\mathrm{OH^-}$ is consumed primarily by $\mathrm{NH_3}$ to form $\mathrm{NH_4^+}$, so the pH decreases slightly.
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.
A student mixes $\mathrm{CH_3NH_2(aq)}$ (methylamine) and $\mathrm{CH_3NH_3Cl(aq)}$ to form a buffer. The student adds a small amount of $\mathrm{HCl(aq)}$. Which statement best explains how the buffer resists a large pH change?
The pH stays exactly constant because buffers neutralize all added strong acid without any limit.
The pH changes little because the buffer is essentially a strong base solution, so $\mathrm{HCl}$ is completely removed.
The added $\mathrm{H^+}$ is consumed primarily by $\mathrm{CH_3NH_3^+}$ to form $\mathrm{CH_3NH_2}$, reducing the increase in $\mathrm{[H^+]}$.
The added $\mathrm{H^+}$ is consumed primarily by $\mathrm{CH_3NH_2}$ to form $\mathrm{CH_3NH_3^+}$, reducing the increase in $\mathrm{[H^+]}$.
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.
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?
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.
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.
Because $\mathrm{HCl}$ is strong, it completely determines the pH after addition, and the $\mathrm{HCOOH/HCOO^-}$ pair does not react significantly.
The added $\mathrm{H^+}$ reacts mainly with $\mathrm{Na^+}$ to form $\mathrm{NaH}$, removing acid and keeping the formic acid equilibrium unchanged.
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.
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.