The ratio of [A⁻] to [HA] for aspirin at pH 7.4 is 10,000 (rounded to the nearest integer). Aspirin, a weak acid, follows the Henderson-Hasselbalch equation where the deprotonated form dominates significantly at this physiological pH.

Problem Solution

Aspirin (HA) dissociates as HA ⇌ H⁺ + A⁻ with Ka = 3.98 × 10-4. First, calculate pKa:

pKa = -log10(Ka) = -log10(3.98 × 10-4) = 3.40

At equilibrium in a buffer at pH 7.4, use the Henderson-Hasselbalch equation:

pH = pKa + log10([A⁻]/[HA])

Rearranging gives:

[A⁻]/[HA] = 10(pH-pKa) = 10(7.4-3.40) = 104.00 = 10,000

This ratio holds regardless of total aspirin concentration, as the equation depends only on pH and pKa for weak acids.

Step-by-Step Derivation

  • Step 1: Compute pKa from given Ka to confirm aspirin’s weak acid nature (pKa ≈ 3.4 matches literature values around 3.5).
  • Step 2: Note pH 7.4 exceeds pKa by 4 units, so A⁻ predominates (ratio > 1).
  • Step 3: Apply 10ΔpH where ΔpH = 7.4 – 3.40 = 4.00, yielding exactly 10,000.
  • Step 4: Round to nearest integer: already 10,000.

No options provided in query, but common MCQ distractors include unrounded values (e.g., 9,997), inverse ratio (0.0001), or pH/pKa confusion—correct is 10,000 via direct calculation.

Aspirin A- HA ratio pH 7.4 is crucial for CSIR NET Life Sciences aspirants studying weak acid equilibria. With Ka of aspirin 3.98 × 10⁻⁴, this aqueous solution problem tests Henderson-Hasselbalch mastery at physiological pH.

Equilibrium Basics

Aspirin (acetylsalicylic acid, HA) ionizes: HA ⇌ H⁺ + A⁻. pKa = 3.40 means it’s mostly protonated below pH 3.4, deprotonated above. At blood pH 7.4, A⁻ form prevails for solubility and pharmacology.

Detailed Calculation

[A⁻]/[HA] = 107.4-3.40 = 10,000

This 10,000:1 ratio explains aspirin’s ionization in plasma.

CSIR NET Relevance

Such problems assess buffer calculations without approximation errors. Practice: vary pH to see ratio shifts (e.g., pH 3.4 = 1).