Properties of Buffers
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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 changes little because $\mathrm{Na^+}$ consumes $\mathrm{OH^-}$ to form $\mathrm{NaOH(aq)}$, removing base from solution.
The pH stays exactly the same because the buffer converts all added base into water with no change in buffer composition.
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 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 pH does not change because $\mathrm{HF}$ is a strong acid and overwhelms the effect of adding $\mathrm{HBr}$.
The pH changes little because $\mathrm{NaF}$ is a strong base that neutralizes all added $\mathrm{HBr}$ completely.
The pH stays exactly constant because $\mathrm{Br^-}$ reacts with water to remove $\mathrm{H^+}$ from solution.
The added $\mathrm{H^+}$ is consumed primarily by $\mathrm{HF}$ to form $\mathrm{F^-}$, which prevents a pH decrease.
The added $\mathrm{H^+}$ is consumed primarily by $\mathrm{F^-}$ to form $\mathrm{HF}$, so the pH decreases only slightly.
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 $\mathrm{H_2CO_3}$ to produce $\mathrm{CO_3^{2-}}$, so the buffer resists pH change.
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 $\mathrm{NO_3^-}$ to form $\mathrm{HNO_3}$, which keeps the pH nearly constant.
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 $\mathrm{Cl^-}$ from $\mathrm{HCl}$ reacts with water to form $\mathrm{OH^-}$, offsetting the added $\mathrm{H^+}$.
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 $\mathrm{NaC_2H_3O_2}$ 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.
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 pH remains constant because buffers eliminate all added $\mathrm{H^+}$ without changing the amounts of buffer components.
The added $\mathrm{H^+}$ reacts mainly with $\mathrm{HClO}$ to form $\mathrm{ClO^-}$, so the pH does not decrease.
The pH changes little because $\mathrm{Cl^-}$ formed from the acid reacts with water to produce $\mathrm{OH^-}$ that neutralizes $\mathrm{H^+}$.
The pH does not change because $\mathrm{NaClO}$ makes the solution strongly basic, preventing any effect from added acid.
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.
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 $\mathrm{NO_2^-}$ 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 added $\mathrm{OH^-}$ is consumed primarily by $\mathrm{HNO_2}$ to form $\mathrm{NO_2^-}$ and $\mathrm{H_2O}$, limiting the pH increase.
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 $\mathrm{NH_3}$ to form $\mathrm{NH_4^+}$, so the pH decreases 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 $\mathrm{NH_4^+}$ to form $\mathrm{NH_3}$ and $\mathrm{H_2O}$, so the pH increases only 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 changes little because $\mathrm{Cl^-}$ reacts with $\mathrm{H^+}$ to form $\mathrm{HCl}$, keeping $\mathrm{[H^+]}$ low.
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 $\mathrm{CH_3NH_2}$ to form $\mathrm{CH_3NH_3^+}$, reducing the increase in $\mathrm{[H^+]}$.
The pH stays exactly constant because buffers neutralize all added strong acid without any limit.
The added $\mathrm{H^+}$ is consumed primarily by $\mathrm{CH_3NH_3^+}$ to form $\mathrm{CH_3NH_2}$, reducing the increase in $\mathrm{[H^+]}$.
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 solution contains methylamine, $\mathrm{CH_3NH_2}$, and methylammonium nitrate, $\mathrm{CH_3NH_3NO_3}$. A small amount of $\mathrm{HCl(aq)}$ is added. Which statement best describes the buffering mechanism and which species reacts with the added acid?
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.
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.
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.
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.
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.
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+.
A buffer contains CH$_3$NH$_2$(aq) and CH$_3$NH$_3$Cl(aq). A small amount of HCl(aq) is added. Which statement best explains the buffer action and identifies which component reacts with the added acid?
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.
The added H$^+$ is consumed mainly by water, and the buffer resists pH change because water is amphoteric.
The added H$^+$ is consumed mainly by CH$_3$NH$_3^+$, which donates OH$^-$ to neutralize the acid and keep pH constant.
The buffer prevents pH change because CH$_3$NH$_2$ is a strong base, so added HCl has no effect on pH.
The added HCl is consumed mainly by Cl$^-$, forming Cl$_2$ and removing H$^+$ so the pH does not change.
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⁺.