MCAT Chemical and Physical Foundations of Biological Systems Quiz: 5a Acid Base Equilibria

What this quiz covers

This quiz focuses on 5a Acid Base Equilibria, giving you a quick way to practice the rules, question types, and explanations that matter most for MCAT Chemical and Physical Foundations of Biological Systems.

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. Acetate (CH$_3$COO$^-$) binds $\mathrm{H^+}$ to form acetic acid, partially resisting the pH drop. (correct answer)
2. Acetic acid dissociates further to generate more $\mathrm{H^+}$, amplifying the pH drop.
3. The added $\mathrm{H^+}$ is consumed by water to form $\mathrm{OH^-}$, so pH rises.
4. No species change occurs because buffers maintain constant pH regardless of added acid.

Explanation: This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically the acetate buffer's response to acid secretion. In the equilibrium CH₃COOH ⇌ H⁺ + CH₃COO⁻, acetate ion (CH₃COO⁻) acts as the conjugate base. When additional H⁺ is secreted into the tubular fluid, acetate binds these protons to form acetic acid (CH₃COOH), shifting the equilibrium to the left. This buffering action partially resists the pH drop that would otherwise occur from the acid secretion. Choice A correctly describes this mechanism. Choice B incorrectly suggests that acetic acid would dissociate further when H⁺ is added, violating Le Châtelier's principle. In renal physiology, various buffer systems help maintain acid-base balance during H⁺ secretion.

Question 2

1. It is suitable because effective buffering typically spans about $pK_a\pm 1$ pH unit. (correct answer)
2. It is unsuitable because buffers only work at exactly $pH=pK_a$.
3. It is unsuitable because buffering is strongest when pH is far from $pK_a$.
4. It is suitable only if benzoic acid is replaced with a strong acid to increase capacity.

Explanation: This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically buffer range and effectiveness. Buffer systems typically maintain effective pH control within approximately pKa ± 1 pH unit, where sufficient amounts of both the acid and conjugate base forms are present. For benzoic acid/benzoate with pKa = 4.20, the effective buffer range spans approximately pH 3.2 to 5.2, which exactly matches the experimental requirement. Choice A correctly identifies this suitable buffer range. Choice B incorrectly limits buffering to exactly pH = pKa, when buffers actually work over a range. When selecting buffers for experiments, ensure the required pH range falls within pKa ± 1 for optimal performance.

Question 3

1. $pK_a\approx 4.8$ because at half-equivalence $\mathrm{pH}=pK_a$. (correct answer)
2. $pK_a\approx 7.0$ because half-equivalence implies neutrality.
3. $pK_a\approx 9.2$ because the conjugate base dominates at half-equivalence.
4. $pK_a$ cannot be inferred from titration data without calculating $K_w$.

Explanation: 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.

Question 4

1. pH decreases because removing $\mathrm{CO_2(aq)}$ shifts equilibrium right, generating more $\mathrm{H^+}$.
2. pH increases because decreasing the acid term increases the ratio $[\mathrm{HCO_3^-}]/[\mathrm{CO_2(aq)}]$. (correct answer)
3. pH remains constant because $[\mathrm{HCO_3^-}]$ is unchanged and dominates the pH.
4. pH becomes exactly 6.1 because that is the $pK_a$ of the bicarbonate system.

Explanation: 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 bicarbonate buffer in response to CO2 removal. Choice B correctly describes pH increase because lower [CO2(aq)] raises the [HCO3-]/[CO2(aq)] ratio. Choice A fails because it predicts more H+ generation, which is a common error when misapplying Le Châtelier to open systems. In future questions, ensure buffer systems are evaluated by considering both capacity and range, especially in open systems with fixed components.

Question 5

1. pH is 7.0 because equal moles of acid and base always yield a neutral solution.
2. pH is less than 7.0 because the solution contains primarily $\mathrm{BH^+}$, a weak acid, at equivalence. (correct answer)
3. pH is greater than 7.0 because the conjugate base B remains in excess at equivalence.
4. pH equals 9.0 because $\mathrm{pH}=pK_a$ at the equivalence point by definition.

Explanation: 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.

Question 6

1. The lactic buffer is highly effective at pH 7.00 because any weak acid buffer works best when pH is far above its $pK_a$.
2. The lactic buffer has limited effectiveness at pH 7.00 because buffering is strongest within about $\pm 1$ pH unit of $pK_a$. (correct answer)
3. The lactic buffer will prevent any pH change until all lactate is converted to lactic acid at an equivalence point.
4. The lactic buffer becomes more effective as pH rises further above $pK_a$ because dissociation is driven to completion.

Explanation: 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 lactic acid buffer at pH far above pKa. Choice B correctly describes limited effectiveness because buffering is strongest within ±1 pH unit of pKa, where both forms are significant. Choice A fails because it claims effectiveness far above pKa, which is a common error when misunderstanding that deprotonated form dominates and cannot buffer acids well. In future questions, ensure buffer systems are evaluated by considering both capacity and range, checking pH proximity to pKa for optimal function.

Question 7

1. pH increases because CO$_2$ production consumes $\mathrm{H^+}$ to form $\mathrm{H_2CO_3}$.
2. pH decreases because increasing $[\mathrm{CO_2(aq)}]$ lowers the ratio $[\mathrm{HCO_3^-}]/[\mathrm{CO_2(aq)}]$. (correct answer)
3. pH remains constant because the bicarbonate concentration is fixed at 25 mM.
4. pH becomes exactly 6.1 because the system is forced to p$K_a$ in a closed container.

Explanation: 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 closed-system bicarbonate buffer with accumulating CO2. Choice B correctly describes pH decrease because higher [CO2(aq)] lowers the [HCO3-]/[CO2(aq)] ratio. Choice A fails because it predicts pH increase from H+ consumption, which is a common error when misinterpreting carbonic acid formation. In future questions, ensure buffer systems are evaluated by considering both capacity and range, distinguishing closed from open systems.

Question 8

1. The pH decreases, and $[\mathrm{HCO_3^-}]$ decreases by approximately $1.0\ \mathrm{mM}$ as it neutralizes added $\mathrm{H^+}$. (correct answer)
2. The pH increases because added $\mathrm{H^+}$ shifts the equilibrium to consume $\mathrm{H^+}$ and generate more $\mathrm{HCO_3^-}$.
3. The pH is unchanged because buffers maintain constant pH until the equivalence point is reached.
4. The pH decreases, and $[\mathrm{HCO_3^-}]$ increases because $\mathrm{CO_2}$ is converted to $\mathrm{HCO_3^-}$ in the presence of acid.

Explanation: This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically the bicarbonate buffer system's response to acid addition. The bicarbonate buffer system maintains pH through the equilibrium CO₂(aq) + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻, where added H⁺ is neutralized by HCO₃⁻ to form H₂CO₃/CO₂. In this passage, when 1.0 mmol of H⁺ is added to 1.0 L containing 24 mM HCO₃⁻, the bicarbonate acts as a base to consume the added acid. Choice A correctly describes that pH decreases (due to added acid) and [HCO₃⁻] decreases by approximately 1.0 mM as it neutralizes the 1.0 mmol of added H⁺. Choice B incorrectly suggests pH increases when acid is added, which violates fundamental acid-base principles. In future questions, remember that buffers resist but don't prevent pH changes, and the conjugate base component consumes added acid.

Question 9

1. pH remains approximately the same because the ratio $[\mathrm{A^-}]/[\mathrm{HA}]$ is unchanged, but buffering capacity decreases due to fewer moles per liter. (correct answer)
2. pH increases because dilution lowers $[\mathrm{H^+}]$ directly, making the solution more basic.
3. pH decreases because dilution shifts the dissociation equilibrium right, generating more $\mathrm{H^+}$.
4. Both pH and buffering capacity are unchanged because dilution does not affect equilibria.

Explanation: This question assesses understanding of acid-base equilibria in buffer solutions (5A). Acid-base equilibria in buffers maintain pH through the ratio of conjugate acid to base, as described by the Henderson-Hasselbalch equation. In this case, the buffer is prepared at pH 7.0 equal to pKa, implying equal concentrations of HA and A- before dilution. Choice A correctly states that pH remains approximately the same because dilution proportionally reduces both [HA] and [A-], preserving their ratio, while capacity decreases due to lower overall concentrations. Choice B is incorrect as it assumes dilution directly affects [H+] independently, overlooking the buffer's role in resisting pH changes, a misconception from treating buffers like strong acids. In future questions, verify buffer behavior by applying the Henderson-Hasselbalch equation before and after changes. Always distinguish between pH stability and buffering capacity when assessing dilutions or additions to buffer systems.

Question 10

1. The equilibrium shifts left, decreasing $[\mathrm{H^+}]$ and increasing pH. (correct answer)
2. The equilibrium shifts right, increasing $[\mathrm{H^+}]$ and decreasing pH.
3. The $K_a$ decreases because adding conjugate base changes the intrinsic acid strength.
4. No shift occurs because weak acids are unaffected by common ions.

Explanation: This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically the common ion effect on weak acid equilibria. When the conjugate base (A⁻) of a weak acid (HA) is added to the solution, it acts as a common ion that shifts the equilibrium HA ⇌ H⁺ + A⁻ to the left according to Le Châtelier's principle. This leftward shift decreases [H⁺], resulting in an increase in pH. Choice A correctly identifies this common ion effect and its consequence on pH. Choice C incorrectly suggests that Ka changes, but equilibrium constants are only affected by temperature, not by concentration changes. When analyzing equilibrium shifts, remember that adding products drives the reaction toward reactants.

Question 11

1. pH increases because decreasing $\mathrm{CO_2}$ shifts the equilibrium left, lowering $[\mathrm{H^+}]$. (correct answer)
2. pH decreases because decreasing $\mathrm{CO_2}$ shifts the equilibrium right, raising $[\mathrm{H^+}]$.
3. pH remains constant because $pK_a$ fixes pH regardless of reactant concentrations.
4. pH decreases because removing an acid removes buffering capacity, forcing $\mathrm{HCO_3^-}$ to dissociate more.

Explanation: This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically the effect of CO₂ changes on blood pH. In the bicarbonate equilibrium CO₂(aq) + H₂O ⇌ H⁺ + HCO₃⁻, decreasing CO₂ concentration shifts the equilibrium to the left according to Le Châtelier's principle. This leftward shift consumes H⁺ ions to form more CO₂ and H₂O, thereby decreasing [H⁺] and increasing pH. Choice A correctly describes this respiratory alkalosis mechanism where hyperventilation removes CO₂ and raises pH. Choice B incorrectly predicts the opposite effect, confusing the direction of equilibrium shift. In respiratory physiology, remember that CO₂ removal (hyperventilation) increases pH while CO₂ retention (hypoventilation) decreases pH.

Question 12

1. Both buffers show the same pH change because they have the same $[\mathrm{A^-}]/[\mathrm{HA}]$ ratio.
2. Buffer X shows a smaller pH change because it has greater buffer capacity (higher total concentration of conjugate pair). (correct answer)
3. Buffer Y shows a smaller pH change because dilute buffers resist pH change more strongly.
4. Neither buffer changes pH because weak acid buffers neutralize all added strong acid until equivalence is reached.

Explanation: This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically the concept of buffer capacity. Buffer capacity depends on the total concentration of the conjugate acid-base pair - higher concentrations provide more molecules to neutralize added acid or base. Both buffers have the same initial pH since they have identical [A⁻]/[HA] ratios, but Buffer X has 10-fold higher total concentration (100 mM vs 10 mM total). Choice B correctly identifies that Buffer X shows a smaller pH change because it has greater buffer capacity due to higher total concentration of the conjugate pair. Choice C incorrectly suggests dilute buffers resist pH change more strongly, when the opposite is true - concentrated buffers have greater capacity. In future problems, remember that buffer capacity increases with total buffer concentration, not just the ratio of components.

Question 13

1. Buffering is most effective when pH is within about $\pm 1$ unit of $pK_a$, where both HBz and Bz$^-$ are present in comparable amounts. (correct answer)
2. Buffering is most effective at the equivalence point because all HBz has been converted to Bz$^-$.
3. Buffering is most effective at very low pH because HBz is maximally protonated and can absorb added acid.
4. Buffering is most effective at very high pH because Bz$^-$ is maximally deprotonated and can absorb added base.

Explanation: This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically the effective buffering range of weak acid-base pairs. Buffer capacity is maximized when both the weak acid and conjugate base are present in significant, comparable amounts, which occurs when pH is near pKa. The Henderson-Hasselbalch equation shows that when pH = pKa ± 1, the ratio [A⁻]/[HA] ranges from 0.1 to 10, providing substantial amounts of both species. Choice A correctly identifies that buffering is most effective when pH is within about ±1 unit of pKa, where both HBz and Bz⁻ are present in comparable amounts. Choice B incorrectly identifies the equivalence point, where essentially all HBz is converted to Bz⁻, leaving no weak acid to neutralize added base. In future buffer problems, remember the practical buffering range is pKa ± 1 pH unit.

Question 14

1. pH will decrease further because lowering $P_{\mathrm{CO_2}}$ drives more $\mathrm{H_2CO_3}$ formation and releases $\mathrm{H^+}$.
2. pH will increase because reducing dissolved $\mathrm{CO_2}$ shifts the equilibrium to consume $\mathrm{H^+}$ and form $\mathrm{CO_2}$ and $\mathrm{H_2O}$. (correct answer)
3. pH will be unchanged because only bicarbonate concentration (not $P_{\mathrm{CO_2}}$) affects the acid–base ratio.
4. pH will increase because $\mathrm{CO_2}$ is a base and its removal decreases basicity of the solution.

Explanation: This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically respiratory compensation in the bicarbonate buffer system. The bicarbonate buffer equilibrium CO2(aq) + H2O ⇌ H2CO3 ⇌ H+ + HCO3- responds to changes in PCO2 according to Le Chatelier's principle. In this passage, after lactic acid addition creates excess H+, hyperventilation reduces PCO2 from 40 to 30 mmHg, which decreases dissolved CO2 concentration. Choice B correctly explains that pH will increase because reducing dissolved CO2 shifts the equilibrium to consume H+ and form CO2 and H2O, partially compensating for the acidosis. Choice A incorrectly suggests lowering PCO2 would decrease pH further, which contradicts the equilibrium response to CO2 removal. In future questions involving respiratory compensation, remember that decreasing PCO2 always increases pH (reduces acidity) as the system consumes H+ to maintain equilibrium.

Question 15

1. Most added $\mathrm{H^+}$ is consumed by $\mathrm{HPO_4^{2-}}$ to form $\mathrm{H_2PO_4^-}$, so pH decreases only modestly. (correct answer)
2. Most added $\mathrm{H^+}$ is consumed by $\mathrm{H_2PO_4^-}$ to form $\mathrm{H_3PO_4}$, so pH increases slightly.
3. The buffer fails because buffering occurs only at $\mathrm{pH} = pK_a \pm 0.1$, so pH changes drastically.
4. Added $\mathrm{H^+}$ causes $\mathrm{H_2PO_4^-}$ to dissociate more, increasing $[\mathrm{HPO_4^{2-}}]$ and raising pH.

Explanation: This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically how phosphate buffers respond to added acid. The phosphate buffer system H2PO4- ⇌ H+ + HPO42- has a pKa of 7.2, and the initial solution contains equal concentrations of both forms, placing the pH at 7.2. In this passage, when 1.0 mmol of strong acid is added to 1.0 L containing 20 mM each of H2PO4- and HPO42-, the added H+ will be consumed by the conjugate base HPO42-. Choice A correctly states that most added H+ is consumed by HPO42- to form H2PO4-, so pH decreases only modestly due to the buffer's capacity. Choice B incorrectly identifies H2PO4- as the species consuming H+, when it's actually the conjugate base HPO42- that acts as the proton acceptor. In future buffer problems, identify which species acts as the base (proton acceptor) - it's always the deprotonated form in the conjugate pair.

Question 16

1. pH increases because lowering $[\mathrm{HCO_3^-}]$ reduces buffering, allowing $[\mathrm{H^+}]$ to fall.
2. pH decreases because the base-to-acid ratio decreases when $[\mathrm{HCO_3^-}]$ is halved at constant CO$_2$. (correct answer)
3. pH is unchanged because $pK_a$ is constant and therefore fixes pH for any mixture.
4. pH increases because CO$_2$ is the conjugate base of $\mathrm{HCO_3^-}$ and remains constant.

Explanation: 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.

Question 17

1. The added $\mathrm{H^+}$ is primarily consumed by Lac$^-$ to form HLac, so the pH decreases only modestly. (correct answer)
2. The added $\mathrm{H^+}$ is primarily consumed by HLac to form Lac$^-$, so the pH increases.
3. The pH change is large because buffering is ineffective when $[\mathrm{HLac}] = [\mathrm{Lac^-}]$.
4. The pH is unchanged because weak acid buffers maintain pH exactly at $pK_a$ regardless of added acid.

Explanation: This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically buffer action when strong acid is added. Buffer systems resist pH changes by converting added H⁺ or OH⁻ to weak acid or conjugate base forms. In this lactate buffer system with equal concentrations of HLac and Lac⁻, added H⁺ ions will react with the conjugate base (Lac⁻) to form more weak acid (HLac), following the equilibrium HLac ⇌ H⁺ + Lac⁻. Choice A correctly states that added H⁺ is consumed by Lac⁻ to form HLac, resulting in only a modest pH decrease due to the buffer's resistance. Choice B incorrectly reverses the reaction direction - HLac cannot consume H⁺ as it would violate the equilibrium principle. In future buffer problems, identify which species (weak acid or conjugate base) will react with the added strong acid or base.

Question 18

1. pH increases because added CO$_2$ consumes $\mathrm{H^+}$ as equilibrium shifts left toward $\mathrm{H_2CO_3}$.
2. pH decreases because increased $[\mathrm{CO_2(aq)}]$ drives formation of $\mathrm{H^+}$, lowering the $[\mathrm{HCO_3^-}]/[\mathrm{CO_2}]$ ratio. (correct answer)
3. pH remains unchanged because buffers completely resist pH change until $[\mathrm{HCO_3^-}]$ is exhausted.
4. pH increases because $[\mathrm{HCO_3^-}]$ rises proportionally with $[\mathrm{CO_2(aq)}]$ in an open system.

Explanation: This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically the bicarbonate buffer system's response to CO₂ changes. The bicarbonate buffer system maintains pH through the equilibrium CO₂(aq) + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻, where increased CO₂ drives the equilibrium rightward. When CO₂(aq) increases by 50%, the system responds by producing more H⁺ ions, which decreases pH according to the Henderson-Hasselbalch equation: pH = pKa + log([HCO₃⁻]/[CO₂]). Choice B correctly identifies that increased [CO₂(aq)] drives H⁺ formation and lowers the [HCO₃⁻]/[CO₂] ratio, resulting in decreased pH. Choice A incorrectly suggests CO₂ consumes H⁺, when actually CO₂ produces H⁺ through carbonic acid formation. In future questions, remember that increased CO₂ always leads to decreased pH in aqueous systems due to carbonic acid formation.

Question 19

1. pH is approximately equal to $pK_a$ because $[\mathrm{HCOOH}] \approx [\mathrm{HCOO^-}]$. (correct answer)
2. pH is 7.0 because the solution is buffered and therefore neutral.
3. pH is greater than 7.0 because NaOH is a strong base and dominates before equivalence.
4. pH is approximately equal to $\tfrac{1}{2}pK_a$ because only half the acid remains.

Explanation: 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.

Question 20

1. pH decreases because converting HA to A$^-$ generates $\mathrm{H^+}$ directly from the weak acid equilibrium.
2. pH increases slightly because base converts HA to A$^-$, increasing the already-large $[\mathrm{A^-}]/[\mathrm{HA}]$ ratio. (correct answer)
3. pH remains constant because buffers prevent any pH change regardless of added base.
4. pH decreases because adding base shifts $\mathrm{HA \rightleftharpoons H^+ + A^-}$ to the right by Le Châtelier's principle.

Explanation: This question assesses understanding of acid-base equilibria in physiological systems (5A), specifically buffer response to base addition when the conjugate base already predominates. Starting with [A⁻] = 10×[HA], the buffer has pH > pKa according to Henderson-Hasselbalch: pH = 5.0 + log(10) = 6.0. When strong base is added, it converts some HA to A⁻, further increasing the already-large [A⁻]/[HA] ratio. Choice B correctly identifies that pH increases slightly because the base converts HA to A⁻, though the change is small due to the already high ratio. Choice D incorrectly invokes Le Châtelier's principle in the wrong direction - adding base consumes H⁺ and shifts equilibrium left, not right. In future buffer problems, remember that adding base always increases pH by increasing the conjugate base to weak acid ratio.