Acid Dissociation Calculators

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Acid dissociation describes the equilibrium process by which an acid releases a proton (H⁺) in aqueous solution. The acid dissociation constant Ka measures the extent of this process: larger Ka means a stronger acid that dissociates more completely. pKa = −log₁₀(Ka) converts Ka to a more convenient scale. Strong acids (Ka >> 1) dissociate essentially completely; weak acids (Ka << 1) dissociate partially. The pKa is used in the Henderson-Hasselbalch equation to calculate pH of buffer solutions and to predict the ionization state of acids and bases at any pH.

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Acid Dissociation Equilibrium

HA ⇌ H⁺ + A⁻

Ka = [H⁺][A⁻] / [HA]

pKa = −log₁₀(Ka)

Lower pKa = stronger acid (dissociates more). Higher pKa = weaker acid.

Ka Examples

  • HCl (strong acid): Ka ≈ 10⁷ (pKa ≈ −7) — effectively complete dissociation
  • Acetic acid (CH₃COOH): Ka = 1.8 × 10⁻⁵ (pKa = 4.75)
  • Carbonic acid (H₂CO₃): Ka₁ = 4.3 × 10⁻⁷ (pKa = 6.37)
  • Ammonium ion (NH₄⁺): Ka = 5.6 × 10⁻¹⁰ (pKa = 9.25)
  • Water (H₂O): Ka = 1.8 × 10⁻¹⁶ (pKa = 15.74)

pH of Weak Acid

For weak acid HA at concentration C: [H⁺] = √(Ka × C) (valid when Ka << C).

pH = ½(pKa − log C).

Example: 0.1 M acetic acid (Ka = 1.8 × 10⁻⁵): [H⁺] = √(1.8 × 10⁻⁵ × 0.1) = 1.34 × 10⁻³ M; pH = 2.87.

Henderson-Hasselbalch Equation

pH = pKa + log([A⁻]/[HA])

At pH = pKa: [A⁻] = [HA] (50% each). Used to design buffers and to calculate the fraction ionized at any pH. In biology: determine ionization state of amino acid side chains, drugs, and metabolites.

Glossary

Ka (Acid Dissociation Constant)
Ka = [H⁺][A⁻]/[HA]; measures the extent of acid dissociation; larger Ka = stronger acid; converted to pKa by: pKa = −log₁₀(Ka).
pKa
−log₁₀(Ka); the pH at which a weak acid is 50% dissociated (equal concentrations of HA and A⁻); lower pKa = stronger acid; used in Henderson-Hasselbalch equation.
Henderson-Hasselbalch Equation
pH = pKa + log([A⁻]/[HA]); calculates buffer pH from the ratio of conjugate base to weak acid; used to design buffers and predict ionization state of molecules at any pH.

Frequently Asked Questions

Ka is the acid dissociation constant: Ka = [H⁺][A⁻]/[HA]. It quantifies how much an acid dissociates in water — larger Ka means more dissociation, stronger acid. pKa = −log₁₀(Ka), converting Ka to a more convenient scale. Strong acids have low (often negative) pKa; weak acids have pKa typically between 2 and 12. Example: acetic acid Ka = 1.8 × 10⁻⁵, pKa = 4.75 — at pH 4.75, exactly half is in the HA form and half is dissociated (A⁻).

For a weak acid HA at concentration C: assume x = [H⁺] at equilibrium. Ka = x²/(C−x) ≈ x²/C (valid when Ka << C). Therefore x = √(Ka × C). pH = −log(x). Example: 0.050 M formic acid (HCOOH), Ka = 1.77 × 10⁻⁴: [H⁺] = √(1.77 × 10⁻⁴ × 0.050) = √(8.85 × 10⁻⁶) = 2.97 × 10⁻³ M; pH = −log(2.97 × 10⁻³) = 2.53. Verify assumption: Ka/C = 0.00354 ≪ 1, so approximation is valid (< 5% error).

pH = pKa + log([A⁻]/[HA]). It calculates the pH of a buffer solution from the ratio of conjugate base to weak acid. At pH = pKa, the ratio = 1 (equal amounts). Buffers are most effective within ±1 pH unit of pKa. Uses: (1) Designing buffers for biochemical experiments (e.g., phosphate buffer for pH 7.4 using pKa₂ = 7.2 of H₂PO₄⁻/HPO₄²⁻). (2) Calculating ionization state of drugs (important for absorption — only un-ionized form crosses membranes). (3) Predicting charge state of amino acid side chains at physiological pH.

Strong acids (HCl, H₂SO₄, HNO₃, HBr, HI, HClO₄) dissociate essentially completely in water — Ka >> 1, pKa < 0. [H⁺] ≈ initial acid concentration. pH = −log(C). Weak acids (acetic acid, carbonic acid, phosphoric acid, most organic acids) dissociate partially — Ka << 1. [H⁺] = √(Ka × C) ≪ C. At the same concentration, a weak acid has higher pH than a strong acid. Degree of dissociation = √(Ka/C) × 100% — increases with dilution. Most biological acids are weak (pKa 2–12), which is why buffers can maintain stable pH.