Acid–Base Equilibria (5A)
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MCAT Chemical and Physical Foundations of Biological Systems › Acid–Base Equilibria (5A)
In a renal physiology model, tubular fluid contains an acetate buffer: $\mathrm{CH_3COOH\rightleftharpoons H^+ + CH_3COO^-}$ with $pK_a=4.76$. If the tubule secretes additional $\mathrm{H^+}$ into the fluid, which statement best describes the immediate chemical response of the buffer pair?
Acetic acid dissociates further to generate more $\mathrm{H^+}$, amplifying the pH drop.
No species change occurs because buffers maintain constant pH regardless of added acid.
The added $\mathrm{H^+}$ is consumed by water to form $\mathrm{OH^-}$, so pH rises.
Acetate (CH$_3$COO$^-$) binds $\mathrm{H^+}$ to form acetic acid, partially resisting the pH drop.
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.
To test buffer range, a lab prepares a 0.10 M buffer of benzoic acid/benzoate with $pK_a=4.20$ at pH 4.20. The experiment requires pH stability between 3.2 and 5.2. Based on the passage, which conclusion about the buffer system is most consistent?
It is unsuitable because buffering is strongest when pH is far from $pK_a$.
It is suitable because effective buffering typically spans about $pK_a\pm 1$ pH unit.
It is suitable only if benzoic acid is replaced with a strong acid to increase capacity.
It is unsuitable because buffers only work at exactly $pH=pK_a$.
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.
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?
$pK_a\approx 9.2$ because the conjugate base dominates at half-equivalence.
$pK_a$ cannot be inferred from titration data without calculating $K_w$.
$pK_a\approx 7.0$ because half-equivalence implies neutrality.
$pK_a\approx 4.8$ because at half-equivalence $\mathrm{pH}=pK_a$.
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.
In a tissue bath, CO$_2$ is rapidly removed by vigorous aeration, decreasing dissolved $\mathrm{CO_2(aq)}$ while $\mathrm{HCO_3^-}$ initially remains near 24 mM. Use $\mathrm{pH}=6.1+\log\left(\frac{HCO_3^-}{CO_2(aq)}\right)$. The concept being tested is how loss of the acid component affects pH.
Which system response is most consistent with this model?
pH decreases because removing $\mathrm{CO_2(aq)}$ shifts equilibrium right, generating more $\mathrm{H^+}$.
pH increases because decreasing the acid term increases the ratio $[\mathrm{HCO_3^-}]/[\mathrm{CO_2(aq)}]$.
pH remains constant because $[\mathrm{HCO_3^-}]$ is unchanged and dominates the pH.
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.
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)?
pH is less than 7.0 because the solution contains primarily $\mathrm{BH^+}$, a weak acid, at equivalence.
pH equals 9.0 because $\mathrm{pH}=pK_a$ at the equivalence point by definition.
pH is greater than 7.0 because the conjugate base B remains in excess at equivalence.
pH is 7.0 because equal moles of acid and base always yield a neutral solution.
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.
A metabolic study monitors lactate accumulation in a closed 1.0 L bioreactor containing a pre-set buffer: 40 mM lactic acid/lactate with $pK_a=3.86$, adjusted initially to pH 7.00 by setting $\mathrm{lactate^-}\gg\mathrm{lactic\ acid}$. The concept being tested is buffer action range and limitations.
As lactate production increases total lactic species while pH remains near 7.00 initially, which statement is most consistent with buffer theory?
The lactic buffer becomes more effective as pH rises further above $pK_a$ because dissociation is driven to completion.
The lactic buffer will prevent any pH change until all lactate is converted to lactic acid at an equivalence point.
The lactic buffer has limited effectiveness at pH 7.00 because buffering is strongest within about $\pm 1$ pH unit of $pK_a$.
The lactic buffer is highly effective at pH 7.00 because any weak acid buffer works best when pH is far above its $pK_a$.
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.
During an ischemia simulation, a solution initially buffered at pH 7.4 with 25 mM $\mathrm{HCO_3^-}$ is sealed (no gas exchange). Over time, CO$_2$ is produced by metabolism and accumulates as dissolved $\mathrm{CO_2(aq)}$. Use $\mathrm{pH}=6.1+\log\left(\frac{HCO_3^-}{CO_2(aq)}\right)$. The concept being tested is closed-system accumulation of the acid component.
Which response is most consistent with this setup as $\mathrm{CO_2(aq)}$ rises?
pH remains constant because the bicarbonate concentration is fixed at 25 mM.
pH becomes exactly 6.1 because the system is forced to p$K_a$ in a closed container.
pH increases because CO$_2$ production consumes $\mathrm{H^+}$ to form $\mathrm{H_2CO_3}$.
pH decreases because increasing $[\mathrm{CO_2(aq)}]$ lowers the ratio $[\mathrm{HCO_3^-}]/[\mathrm{CO_2(aq)}]$.
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.
In an ex vivo study of blood acid–base balance, a researcher prepares 1.0 L of a bicarbonate buffer mimicking plasma: $\mathrm{HCO_3^-} = 24\ \mathrm{mM}$ and dissolved $\mathrm{CO_2(aq)} = 1.2\ \mathrm{mM}$ at 37°C. The relevant equilibrium is $\mathrm{CO_2(aq) + H_2O \rightleftharpoons H_2CO_3 \rightleftharpoons H^+ + HCO_3^-}$, and for the overall pair $\mathrm{H_2CO_3/HCO_3^-}$, $pK_a = 6.1$. The system is open to a controlled gas phase such that $\mathrm{CO_2(aq)}$ remains effectively constant during short manipulations. A bolus of strong acid adds $1.0\ \mathrm{mmol}$ of $\mathrm{H^+}$ to the solution, with negligible volume change. Which statement best describes the system’s response?
(Assume $\mathrm{H^+}$ is consumed primarily by $\mathrm{HCO_3^-}$ to form $\mathrm{H_2CO_3/CO_2}$.)
The pH decreases, and $[\mathrm{HCO_3^-}]$ decreases by approximately $1.0\ \mathrm{mM}$ as it neutralizes added $\mathrm{H^+}$.
The pH increases because added $\mathrm{H^+}$ shifts the equilibrium to consume $\mathrm{H^+}$ and generate more $\mathrm{HCO_3^-}$.
The pH decreases, and $[\mathrm{HCO_3^-}]$ increases because $\mathrm{CO_2}$ is converted to $\mathrm{HCO_3^-}$ in the presence of acid.
The pH is unchanged because buffers maintain constant pH until the equivalence point is reached.
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.
A 1.0 L solution contains a buffer pair $\mathrm{HA/A^-}$ with $pK_a=7.0$. It is prepared at pH 7.0. The solution is then diluted tenfold with pure water (to 10.0 L) without adding any acid or base. The concept being tested is how dilution affects buffer pH versus buffer capacity.
Which statement best describes the effect of dilution on pH and buffering capacity?
Both pH and buffering capacity are unchanged because dilution does not affect equilibria.
pH increases because dilution lowers $[\mathrm{H^+}]$ directly, making the solution more basic.
pH remains approximately the same because the ratio $[\mathrm{A^-}]/[\mathrm{HA}]$ is unchanged, but buffering capacity decreases due to fewer moles per liter.
pH decreases because dilution shifts the dissociation equilibrium right, generating more $\mathrm{H^+}$.
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.
A weak acid drug, HA, is dissolved in water at 25°C to a formal concentration of 0.10 M. Its $K_a=1.0\times10^{-5}$. Without doing detailed calculations, which change would most likely occur following addition of a small amount of NaA (the conjugate base) while keeping volume constant?
The equilibrium shifts right, increasing $[\mathrm{H^+}]$ and decreasing pH.
No shift occurs because weak acids are unaffected by common ions.
The $K_a$ decreases because adding conjugate base changes the intrinsic acid strength.
The equilibrium shifts left, decreasing $[\mathrm{H^+}]$ and increasing pH.
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.