Acid–Base Equilibria (5A) - Chem & Physics

A lab prepares a 1.0 L buffer for cell culture using $\mathrm{NH_4^+/NH_3}$ with $pK_a=9.25$ . The buffer is adjusted to pH 7.25. The concept being tested is buffer composition far from $pK_a$ . Which statement best describes the relative amounts of $\mathrm{NH_4^+}$ and $\mathrm{NH_3}$ at pH 7.25?
Answer: $\mathrm{NH_4^+}$ is much greater than $\mathrm{NH_3}$ because pH is well below $pK_a$ . This question assesses understanding of acid-base equilibria in physiological systems (5A). Acid-base equilibria involve the balance between acids and bases in solution, often maintained by buffer systems. In this passage, the system's pH regulation is demonstrated through ammonium buffer composition at pH below pKa. Choice B correctly describes NH4+ >> NH3 because pH well below pKa favors protonated form. Choice A fails because it claims NH3 predominance, which is a common error when inverting pH-pKa for acids vs. bases. In future questions, ensure buffer systems are evaluated by considering both capacity and range, calculating ratios for speciation.
A researcher needs a buffer at pH 8.0 and can choose between two monoprotic weak acids: Acid X with $pK_a=4.0$ and Acid Y with $pK_a=8.0$ . Both can be prepared at the same total concentration. Based on the passage, which conclusion is most consistent with buffer theory?
Answer: Acid Y is preferred because its $pK_a$ is closest to the target pH, maximizing buffering near pH 8.0. This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically buffer selection based on pKa. The Henderson-Hasselbalch equation shows that buffers work most effectively when pH is within approximately ±1 unit of the buffer's pKa. For a target pH of 8.0, Acid Y with pKa = 8.0 is ideal because pH = pKa, allowing equal concentrations of acid and conjugate base for maximum buffer capacity. Choice A correctly identifies this principle of matching pKa to target pH. Choice B incorrectly assumes lower pKa improves buffering at any pH, when actually a pKa of 4.0 would require a very high [A⁻]/[HA] ratio at pH 8.0, reducing buffer effectiveness. Always select buffers with pKa close to the desired working pH.
A blood-gas analyzer is calibrated using a bicarbonate buffer where $pK_a = 6.1$ for the $\mathrm{H_2CO_3}/\mathrm{HCO_3^-}$ pair. Two calibration solutions are prepared at 37°C: Solution 1: $[\mathrm{HCO_3^-}] = 20\ \mathrm{mM}$ , $[\mathrm{CO_2(aq)}] = 1.0\ \mathrm{mM}$ Solution 2: $[\mathrm{HCO_3^-}] = 10\ \mathrm{mM}$ , $[\mathrm{CO_2(aq)}] = 1.0\ \mathrm{mM}$ Assume $[\mathrm{H_2CO_3}]$ tracks with $[\mathrm{CO_2(aq)}]$ and the Henderson–Hasselbalch relationship applies qualitatively. Which statement best describes the system's response when moving from Solution 1 to Solution 2?
Answer: pH decreases because the base-to-acid ratio decreases when $[\mathrm{HCO_3^-}]$ is halved at constant CO $_2$ . This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically the bicarbonate buffer system's pH dependence on component ratios. The Henderson-Hasselbalch equation for this system is: pH = pKa + log([HCO₃⁻]/[CO₂]), where pKa = 6.1. In Solution 1, the ratio is 20/1 = 20, while in Solution 2, it's 10/1 = 10, representing a decrease in the base-to-acid ratio. Choice B correctly identifies that pH decreases because the base-to-acid ratio decreases when [HCO₃⁻] is halved at constant CO₂, resulting in a smaller log term in the Henderson-Hasselbalch equation. Choice A incorrectly suggests that lowering [HCO₃⁻] would cause pH to increase, which contradicts the mathematical relationship. In future problems, remember that decreasing the numerator in the Henderson-Hasselbalch ratio always decreases pH.
A 100 mL solution contains a weak acid/conjugate base pair at 25°C with $pK_a=6.5$ . The solution is initially prepared so that $\mathrm{pH}=5.5$ . The concept being tested is interpreting pH relative to $pK_a$ using Henderson–Hasselbalch. Which statement best describes the ratio $[\mathrm{A^-}]/[\mathrm{HA}]$ at pH 5.5?
Answer: $[\mathrm{A^-}]/[\mathrm{HA}]<1$ because pH is below $pK_a$ . This question assesses understanding of acid-base equilibria in physiological systems (5A). Acid-base equilibria involve the balance between acids and bases in solution, often maintained by buffer systems. In this passage, the system's pH regulation is demonstrated through conjugate ratio at pH below pKa. Choice B correctly describes [A-]/[HA] <1 because pH < pKa favors HA. Choice A fails because it claims ratio >1, which is a common error when reversing Henderson-Hasselbalch implications. In future questions, ensure buffer systems are evaluated by considering both capacity and range, using log ratios for pH-pKa differences.
A weak acid equilibrium is studied in a microfluidic device at constant temperature: $\mathrm{HA\rightleftharpoons H^+ + A^-}$ with fixed $K_a$ . The device removes A $^-$ continuously through a selective membrane while leaving HA and $\mathrm{H^+}$ behind. What change would most likely occur in the remaining solution?
Answer: More HA dissociates to replace removed A $^-$ , increasing $[\mathrm{H^+}]$ and lowering pH. This question assesses understanding of acid-base equilibria and Le Chatelier's principle in physiological systems (5A). When a system at equilibrium is disturbed, it shifts to counteract the disturbance and re-establish equilibrium according to Le Chatelier's principle. In this microfluidic device, the continuous removal of A⁻ ions creates a stress on the equilibrium system. According to Le Chatelier's principle, the equilibrium will shift to the right (forward direction) to replace the removed A⁻, causing more HA to dissociate into H⁺ and A⁻. This increased dissociation produces more H⁺ ions, which increases [H⁺] and consequently lowers the pH of the solution. Choice C is incorrect because removing products does affect equilibrium position - this is a fundamental principle of chemical equilibrium. In future questions involving equilibrium disturbances, always apply Le Chatelier's principle: when products are removed, the equilibrium shifts forward to produce more products, and when reactants are removed, the equilibrium shifts backward to produce more reactants.
To compare two candidate buffers for a perfusion medium, a lab prepares Buffer X: 50 mM acetic acid/acetate with $pK_a=4.76$ , adjusted so $[\mathrm{A^-}]=[\mathrm{HA}]$ ; and Buffer Y: 50 mM Tris/TrisH $^+$ with $pK_a=8.1$ , also adjusted so $[\mathrm{base}]=[\mathrm{acid}]$ . Both are at pH 7.40 initially and at the same temperature. The concept being tested is buffer effectiveness as a function of proximity of pH to $pK_a$ . Which conclusion about resistance to a small addition of strong acid at pH 7.40 is most consistent with acid–base equilibria?
Answer: Buffer Y resists better because its $pK_a$ is closer to 7.40, so the conjugate pair can absorb added $\mathrm{H^+}$ with smaller pH change. This question assesses understanding of acid-base equilibria in physiological systems (5A). Acid-base equilibria involve the balance between acids and bases in solution, often maintained by buffer systems. In this passage, the system's pH regulation is demonstrated through comparing buffer effectiveness based on pKa proximity to target pH. Choice B correctly describes Buffer Y's better resistance because its pKa is closer to 7.40, allowing better absorption of H+ with smaller pH change despite equal concentrations. Choice A fails because it prioritizes concentration over pKa match, which is a common error when overlooking that buffering is maximal near pKa. In future questions, ensure buffer systems are evaluated by considering both capacity and range, prioritizing pKa alignment for effectiveness.
An in vitro blood substitute uses the bicarbonate system with $pK_a=6.10$ . A sample initially has $[\mathrm{HCO_3^-}]=24\ \mathrm{mM}$ and $[\mathrm{CO_2(aq)}]=1.2\ \mathrm{mM}$ . A metabolic inhibitor decreases bicarbonate to 12 mM while CO $_2$ (aq) remains 1.2 mM. What change would most likely occur following this condition?
Answer: pH decreases because the base/acid ratio decreases. This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically the bicarbonate buffer system's response to concentration changes. The Henderson-Hasselbalch equation pH = pKa + log([HCO₃⁻]/[CO₂]) shows that pH depends on the ratio of bicarbonate to dissolved CO₂. When [HCO₃⁻] decreases from 24 mM to 12 mM while [CO₂] remains at 1.2 mM, the base/acid ratio decreases from 20:1 to 10:1. This decrease in the ratio results in a lower pH according to the equation. Choice A correctly identifies this relationship. Choice B incorrectly suggests that removing base would shift equilibrium to consume H⁺, when actually the decreased ratio produces more H⁺. In clinical contexts, decreased bicarbonate with constant CO₂ indicates metabolic acidosis.
A simplified titration curve is recorded for 0.10 M HA (25.0 mL) titrated with 0.10 M NaOH. The measured pH at 12.5 mL added is 4.8. The concept being tested is identifying $pK_a$ from half-equivalence data. Assuming ideal behavior and a monoprotic acid, which conclusion is most consistent with the data?
Answer: $pK_a\approx 4.8$ because at half-equivalence $\mathrm{pH}=pK_a$ . This question assesses understanding of acid-base equilibria in physiological systems (5A). Acid-base equilibria involve the balance between acids and bases in solution, often maintained by buffer systems. In this passage, the system's pH regulation is demonstrated through titration data at half-equivalence. Choice A correctly describes pKa ≈4.8 because pH = pKa at half-equivalence for monoprotic acids. Choice B fails because it assumes neutrality at 7.0, which is a common error when ignoring pKa dependence. In future questions, ensure buffer systems are evaluated by considering both capacity and range, using half-equivalence for pKa estimation.
A researcher tests a weak acid HA in water at 25°C with $pK_a=4.0$ . Two separate 50 mL samples are prepared at the same initial pH using appropriate amounts of HA and NaA: Sample 1 has total buffer concentration 5 mM; Sample 2 has total buffer concentration 50 mM. The concept being tested is buffer capacity under identical initial pH. When 0.10 mmol of HCl is added to each sample, which outcome is most consistent with buffer behavior?
Answer: Sample 2 experiences a smaller pH change because it contains more moles of conjugate base to neutralize added acid. This question assesses understanding of acid-base equilibria in physiological systems (5A). Acid-base equilibria involve the balance between acids and bases in solution, often maintained by buffer systems. In this passage, the system's pH regulation is demonstrated through buffer capacity comparison at same pH but different concentrations. Choice B correctly describes smaller pH change in Sample 2 because higher concentration provides more moles to neutralize acid. Choice C fails because it assumes identical changes from same initial pH, which is a common error when overlooking concentration's role in capacity. In future questions, ensure buffer systems are evaluated by considering both capacity and range, factoring total moles for resistance.
A 10.0 mL sample of 0.10 M weak base B (with conjugate acid $\mathrm{BH^+}$ , $pK_a(\mathrm{BH^+})=9.0$ ) is titrated with 0.10 M HCl at 25°C. The concept being tested is identifying the buffer region and equivalence behavior for weak base titration. Which statement is most consistent with the pH at the equivalence point (after adding 10.0 mL of HCl)?
Answer: pH is less than 7.0 because the solution contains primarily $\mathrm{BH^+}$ , a weak acid, at equivalence. This question assesses understanding of acid-base equilibria in physiological systems (5A). Acid-base equilibria involve the balance between acids and bases in solution, often maintained by buffer systems. In this passage, the system's pH regulation is demonstrated through weak base titration to equivalence. Choice B correctly describes pH < 7.0 because equivalence yields BH+, a weak acid, hydrolyzing to acidic pH. Choice C fails because it assumes excess base, which is a common error when misidentifying equivalence. In future questions, ensure buffer systems are evaluated by considering both capacity and range, noting weak base equivalence is acidic.
A pharmacology team studies a weak acid drug (HA) in aqueous solution at 25°C. The drug has $pK_a = 5.0$ and is formulated at $C_T = [\mathrm{HA}] + [\mathrm{A^-}] = 10\ \mathrm{mM}$ . The solution pH is adjusted from 4.0 to 6.0 using small amounts of strong acid/base, without changing $C_T$ appreciably. Which statement best describes the system's response to raising the pH from 4.0 to 6.0? (Assume only the HA/A $^-$ equilibrium is relevant.)
Answer: The fraction of A $^-$ increases because lower $[\mathrm{H^+}]$ shifts $\mathrm{HA \rightleftharpoons H^+ + A^-}$ to the right. This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically the Henderson-Hasselbalch relationship for weak acid speciation. The equilibrium HA ⇌ H⁺ + A⁻ shifts based on pH relative to pKa, following the Henderson-Hasselbalch equation: pH = pKa + log([A⁻]/[HA]). In this passage, when pH increases from 4.0 to 6.0 (crossing above pKa = 5.0), the lower [H⁺] concentration drives the equilibrium to the right. Choice B correctly describes that the fraction of A⁻ increases because lower [H⁺] shifts the equilibrium rightward, converting HA to A⁻. Choice C incorrectly suggests that adding base consumes A⁻, when actually bases consume acids (HA) to form more conjugate base (A⁻). In future questions, remember that at pH > pKa, the deprotonated form predominates, while at pH < pKa, the protonated form predominates.
To illustrate titration behavior in a biologically relevant context, a researcher titrates 25.0 mL of 0.10 M formic acid (HCOOH) with 0.10 M NaOH at 25°C. For $\mathrm{HCOOH} \rightleftharpoons \mathrm{H^+} + \mathrm{HCOO^-}$ , $K_a = 1.8 \times 10^{-4}$ (so $pK_a \approx 3.74$ ). Assume volumes are additive. Which statement best describes the system's response at the half-equivalence point?
Answer: pH is approximately equal to $pK_a$ because $[\mathrm{HCOOH}] \approx [\mathrm{HCOO^-}]$ . This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically the pH at the half-equivalence point of a weak acid titration. At the half-equivalence point, exactly half of the weak acid has been neutralized by the strong base, creating equal concentrations of the weak acid (HCOOH) and its conjugate base (HCOO⁻). According to the Henderson-Hasselbalch equation: pH = pKa + log([A⁻]/[HA]), when [A⁻] = [HA], the log term equals zero, making pH = pKa. Choice A correctly identifies that pH ≈ pKa because [HCOOH] ≈ [HCOO⁻] at the half-equivalence point. Choice D incorrectly suggests pH = ½pKa, which has no theoretical basis in acid-base chemistry. In future titration problems, remember that the half-equivalence point always occurs when pH = pKa for monoprotic weak acids.
A weak acid drug (HA) is studied in aqueous solution at 25°C: $\mathrm{HA \rightleftharpoons H^+ + A^-}$ with $K_a=1.0\times10^{-5}$ (so $pK_a=5.0$ ). A researcher prepares two 100 mL solutions, each containing 10 mM HA. Solution 2 also contains 10 mM NaA (the conjugate base). The concept being tested is the common-ion effect on weak-acid dissociation. Which statement best describes the relative percent dissociation of HA in Solution 2 compared with Solution 1?
Answer: Lower in Solution 2 because added $\mathrm{A^-}$ shifts equilibrium left, suppressing HA dissociation. This question assesses understanding of acid-base equilibria in physiological systems (5A). Acid-base equilibria involve the balance between acids and bases in solution, often maintained by buffer systems. In this passage, the system's pH regulation is demonstrated through the common-ion effect in weak acid solutions. Choice B correctly describes lower dissociation in Solution 2 because added A- shifts equilibrium left, suppressing HA dissociation per Le Châtelier. Choice A fails because it claims higher dissociation, which is a common error when reversing the common-ion effect. In future questions, ensure buffer systems are evaluated by considering both capacity and range, applying Le Châtelier to added conjugates.
A clinician-researcher models plasma buffering with bicarbonate ( $pK_a=6.10$ ). Two samples have the same $[\mathrm{HCO_3^-}]$ but different dissolved CO $_2$ : Sample 1 has higher CO $_2$ than Sample 2. Based on the buffer equilibrium, which statement best describes their relative pH values?
Answer: Sample 1 has lower pH because higher CO $_2$ corresponds to a lower base/acid ratio. This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically the bicarbonate buffer system's pH dependence on component ratios. According to the Henderson-Hasselbalch equation pH = pKa + log([HCO₃⁻]/[CO₂]), pH depends on the ratio of bicarbonate to dissolved CO₂. With constant [HCO₃⁻], Sample 1's higher [CO₂] creates a lower base/acid ratio compared to Sample 2. This lower ratio results in a lower pH for Sample 1. Choice A correctly identifies this inverse relationship between CO₂ concentration and pH. Choice C incorrectly ignores the role of CO₂ in determining pH, focusing only on bicarbonate. In clinical assessment, both components of the bicarbonate buffer must be considered.
A researcher prepares 100 mL of a buffer intended to mimic intracellular phosphate buffering: $\mathrm{H_2PO_4^- \rightleftharpoons H^+ + HPO_4^{2-}}$ with $pK_a=7.2$ . The solution contains $[\mathrm{H_2PO_4^-}]=30\ \mathrm{mM}$ and $[\mathrm{HPO_4^{2-}}]=30\ \mathrm{mM}$ at 25°C. The concept being tested is buffer capacity near $pK_a$ . If 1.0 mL of 1.0 M NaOH is added and mixed (assume negligible volume change), what change would most likely occur immediately?
Answer: pH increases because $\mathrm{OH^-}$ converts $\mathrm{H_2PO_4^-}$ to $\mathrm{HPO_4^{2-}}$ , increasing the base/acid ratio. This question assesses understanding of acid-base equilibria in physiological systems (5A). Acid-base equilibria involve the balance between acids and bases in solution, often maintained by buffer systems. In this passage, the system's pH regulation is demonstrated through the phosphate buffer responding to added NaOH. Choice A correctly describes the system's response because OH- reacts with H2PO4- to form HPO4(2-), increasing the base/acid ratio and thus pH. Choice C fails because it assumes pH remains at pKa with equal concentrations, which is a common error when ignoring the effect of added base on the ratio. In future questions, ensure buffer systems are evaluated by considering both capacity and range, especially near the pKa where buffering is maximal.
In an ex vivo blood-mimic experiment (37°C), a researcher prepares 1.0 L of a bicarbonate buffer containing $[\mathrm{HCO_3^-}]=24\ \mathrm{mM}$ and dissolved $\mathrm{CO_2(aq)}=1.2\ \mathrm{mM}$ . The relevant equilibrium is $\mathrm{CO_2(aq)+H_2O \rightleftharpoons H_2CO_3 \rightleftharpoons H^+ + HCO_3^-}$ , and for this system the effective $pK_a=6.1$ when $[\mathrm{CO_2(aq)}]$ is used as the acid term. The sample is open to a controlled gas phase that maintains $[\mathrm{CO_2(aq)}]$ constant over the next minute. The concept being tested is buffer response to a small strong-acid load. A bolus of $0.50\ \mathrm{mmol}$ HCl is rapidly mixed into the 1.0 L sample. Based on acid–base equilibria, which conclusion about the immediate pH change is most consistent with the buffer system described?
Answer: pH decreases slightly because added $\mathrm{H^+}$ is largely consumed by $\mathrm{HCO_3^-}$ to form $\mathrm{CO_2(aq)}$ , reducing the base/acid ratio. This question assesses understanding of acid-base equilibria in physiological systems (5A). Acid-base equilibria involve the balance between acids and bases in solution, often maintained by buffer systems. In this passage, the system's pH regulation is demonstrated through the bicarbonate buffer in an open system where [CO2(aq)] is held constant. Choice A correctly describes the system's response to added HCl because the H+ is consumed by HCO3- to form CO2(aq), which is vented to maintain constant [CO2(aq)], slightly reducing the [HCO3-]/[CO2(aq)] ratio and thus pH. Choice C fails because it assumes buffers keep pH exactly constant, which is a common error when overlooking that buffers minimize but do not eliminate pH changes. In future questions, ensure buffer systems are evaluated by considering both capacity and range, especially in open versus closed systems.
A weak acid equilibrium is studied: $\mathrm{HA \rightleftharpoons H^+ + A^-}$ with $pK_a=6.0$ . A 200 mL solution is prepared with $[\mathrm{HA}]=[\mathrm{A^-}]$ (a buffer at pH 6.0). The concept being tested is Le Chtelier's principle in buffer systems. If a small amount of strong acid is added, which shift is most consistent with maintaining equilibrium?
Answer: Shift left, consuming $\mathrm{H^+}$ as $\mathrm{A^-}$ is protonated to HA. This question assesses understanding of acid-base equilibria in physiological systems (5A). Acid-base equilibria involve the balance between acids and bases in solution, often maintained by buffer systems. In this passage, the system's pH regulation is demonstrated through weak acid buffer response to added strong acid. Choice B correctly describes left shift, consuming H+ by protonating A- to HA. Choice A fails because it predicts right shift, which is a common error when confusing stress direction in Le Châtelier. In future questions, ensure buffer systems are evaluated by considering both capacity and range, applying Le Châtelier to added H+.
A researcher titrates 20.0 mL of 0.050 M diprotic acid H $_2$ A but observes two well-separated buffer regions, suggesting $pK_{a1}=3.0$ and $pK_{a2}=8.0$ . The concept being tested is identifying which conjugate pair buffers near a given pH. Near pH 8.0, which conjugate acid–base pair most directly provides buffering capacity?
Answer: $\mathrm{HA^-/A^{2-}}$ because buffering is strongest near $pK_{a2}$ for the second deprotonation. This question assesses understanding of acid-base equilibria in polyprotic systems (5A). Acid-base equilibria involve the stepwise dissociation of acids and the buffering capacity provided by conjugate pairs near their pKa values. In this scenario, the diprotic acid H2A has pKa1 = 3.0 and pKa2 = 8.0, creating two distinct buffering regions during titration. Choice B correctly identifies HA-/A2- as the pair providing buffering near pH 8.0 because this region corresponds to the second deprotonation where buffering is maximal around pKa2. Choice A is incorrect because it focuses on the first deprotonation, which buffers near pH 3.0, a common error in confusing the order of pKa values in polyprotic acids. In future questions, evaluate buffering pairs by matching the target pH to the closest pKa and identifying the corresponding conjugate species. Remember that effective buffering occurs within ±1 pH unit of the pKa for the relevant dissociation step.
A weak acid preservative (HA) in an IV solution follows $\mathrm{HA \rightleftharpoons H^+ + A^-}$ with $pK_a = 4.0$ . The IV solution is formulated at pH 4.0 using a mixture of HA and its conjugate base $\mathrm{A^-}$ (assume no other acid–base species contribute). During storage, a small amount of NaOH contaminant is introduced, adding $\mathrm{OH^-}$ that reacts to completion with available protons. Which statement best describes the immediate effect of the NaOH contamination on the HA/Aaf equilibrium?
Answer: The equilibrium shifts toward Aaf because $\mathrm{OH^-}$ consumes $\mathrm{H^+}$ , promoting further dissociation of HA to restore $[\mathrm{H^+}]$ . This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically Le Chatelier's principle applied to weak acid equilibria when base is added. For the equilibrium HA ⇌ H+ + A-, adding OH- removes H+ through the reaction OH- + H+ → H2O. In this passage, when NaOH contaminant is introduced, the OH- consumes available H+, decreasing [H+] in the system. Choice C correctly explains that the equilibrium shifts toward A- because OH- consumes H+, promoting further dissociation of HA to restore [H+] according to Le Chatelier's principle. Choice A incorrectly suggests the equilibrium shifts toward HA, which would occur if H+ were added, not removed. In future contamination problems, remember that removing a product (H+) drives the equilibrium forward to produce more of that product.
An enzyme assay uses Tris buffer, modeled as $\mathrm{BH^+ \rightleftharpoons H^+ + B}$ with $pK_a = 8.1$ at the assay temperature. The buffer is prepared at total concentration $[\mathrm{BH^+}]+[\mathrm{B}] = 50\ \mathrm{mM}$ and adjusted to pH 8.1. During the reaction, the enzyme produces 2.0 mmol of $\mathrm{H^+}$ in 1.0 L. Which statement best describes the expected change in buffer composition and pH?
Answer: pH decreases modestly and $[\mathrm{B}]$ decreases because B consumes $\mathrm{H^+}$ to form $\mathrm{BH^+}$ . This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically how buffers respond to acid production at pH = pKa. The Tris buffer system BH+ ⇌ H+ + B has pKa = 8.1, and at pH 8.1, equal amounts of BH+ and B are present (25 mM each). In this passage, when 2.0 mmol of H+ is produced in 1.0 L, the conjugate base B will consume this H+ to form BH+. Choice B correctly states that pH decreases modestly and [B] decreases because B consumes H+ to form BH+, with the buffer minimizing but not preventing pH change. Choice C incorrectly claims pH is unchanged, misunderstanding that buffers minimize but don't eliminate pH changes even at optimal pH = pKa. In future enzyme assay problems, remember that buffers work best at pH = pKa but still show measurable pH changes when acid or base is added.

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