General Science20 Concepts & Facts

What Is pKa? Acid Dissociation Constant & Henderson-Hasselbalch Equation

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In physical chemistry, acid strength reflects the thermodynamic tendency of a chemical species to donate a proton to an aqueous solvent. When a generic Bronsted-Lowry acid dissolves in water, it establishes a dynamic equilibrium represented by the reaction HA plus H2O producing hydronium ions and the conjugate base A minus. The equilibrium constant for this reversible dissociation is termed the acid dissociation constant, or Ka. Because Ka values for acids span many orders of magnitude—ranging from over ten raised to the sixth power for strong mineral acids to less than ten raised to the negative fourteenth power for extremely weak acids—chemists express acidity using the logarithmic index pKa. Defined mathematically as the negative base-ten logarithm of Ka, pKa provides a standardized metric where smaller or negative numerical values indicate stronger acids that dissociate more readily.

The molecular structure of an acid directly dictates its pKa through electronic and steric factors. Polyprotic acids, which contain more than one ionizable hydrogen atom, exhibit successive dissociation constants where Ka1 exceeds Ka2, which in turn exceeds Ka3. For example, phosphoric acid exhibits a first pKa of 2.15, a second pKa of 7.20, and a third pKa of 12.35. This sequential increase occurs because removing a positively charged proton from an increasingly negative anion demands greater electrostatic work. In addition, electronegative substituents exert an electron-withdrawing inductive effect that stabilizes the conjugate base anion, substantially reducing pKa. Trichloroacetic acid, with three electronegative chlorine atoms pulling electron density away from the carboxylate group, has a pKa of 0.65, whereas unsubstituted acetic acid has a pKa of 4.76. Resonance stabilization similarly enhances acidity, as seen in phenol, whose conjugate phenoxide anion delocalizes negative charge across an aromatic ring.

Quantitative calculations involving weak acids and buffer solutions rely on the Henderson-Hasselbalch equation. Derived directly by taking the negative logarithm of the Ka expression, this formula defines pH as the sum of pKa and the base-ten logarithm of the ratio of conjugate base concentration to weak acid concentration. When the molar concentrations of the weak acid and its conjugate base are identical, the logarithmic term equals zero, making the solution pH equal to the pKa. This point of equimolar equivalence marks maximum buffering capacity, meaning the solution resists changes in pH most effectively when challenged with small additions of strong acids or bases. In human physiology, this mathematical relationship governs the carbonic acid-bicarbonate buffer system that maintains arterial blood pH within the narrow window of 7.35 to 7.45.

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Ka (acid dissociation constant) is the equilibrium constant quantifying the extent of dissociation of an acid HA in aqueous solution: Ka = [H3O+][A-] / [HA].

pKa is defined mathematically as the negative base-10 logarithm of Ka: pKa = -log10(Ka).

Lower or negative pKa values denote stronger acids with greater degrees of ionization, whereas higher pKa values represent weaker acids.

While pH measures the concentration of hydronium ions in a specific solution, pKa is an intrinsic thermodynamic property of a chemical species at a given temperature.

Strong mineral acids such as hydrochloric acid (HCl, pKa ≈ -6.3) and sulfuric acid (H2SO4, pKa1 ≈ -3.0) dissociate completely in dilute aqueous solution.

Acetic acid (CH3COOH) is a typical weak carboxylic acid with a Ka of 1.74 × 10^-5 at 25 °C, corresponding to a pKa of approximately 4.76.

The Henderson-Hasselbalch equation relates pH, pKa, and conjugate base/acid ratio: pH = pKa + log10([A-] / [HA]).

When the molar concentration of an unprotonated conjugate base equals that of the weak acid ([A-] = [HA]), the logarithmic ratio is zero and pH equals pKa.

The optimal buffering capacity of a conjugate acid-base pair occurs at pH = pKa, with effective buffering maintained within the range pH = pKa ± 1.

#10

Polyprotic acids dissociate in stepwise equilibria, where successive dissociation constants decrease systematically: Ka1 >> Ka2 >> Ka3.

#11

Phosphoric acid (H3PO4) possesses three distinct pKa values: pKa1 = 2.15, pKa2 = 7.20, and pKa3 = 12.35, reflecting increasing electrostatic retention of protons.

#12

Electronegative substituents exert an electron-withdrawing inductive effect (-I), stabilizing the conjugate base and dramatically lowering pKa (e.g., trichloroacetic acid pKa = 0.65 vs acetic acid pKa = 4.76).

#13

Resonance delocalization stabilizes conjugate bases, explaining why phenol (pKa ≈ 9.95) is substantially more acidic than cyclohexanol (pKa ≈ 16).

#14

In titration curves of weak acids with strong bases, the half-equivalence point represents the condition where exactly half the acid is neutralized, establishing pH = pKa.

#15

The carbonic acid-bicarbonate buffer system (H2CO3 / HCO3-) regulates human arterial blood pH within the narrow physiological range of 7.35 to 7.45.

#16

Although the carbonic acid pKa is 6.10, the blood maintains pH 7.4 because the ratio of [HCO3-] to dissolved CO2 is held at approximately 20:1 through pulmonary ventilation.

#17

Amino acids exist as zwitterions with multiple pKa values corresponding to alpha-carboxyl groups (pKa ≈ 2.0-2.5) and alpha-amino groups (pKa ≈ 9.0-10.5).

#18

The isoelectric point (pI) of a neutral diprotic amino acid is calculated as the arithmetic mean of its two pKa values: pI = (pKa1 + pKa2) / 2.

#19

Temperature alters equilibrium constants: acid dissociation is generally endothermic, causing Ka to increase and pKa to decrease slightly as temperature rises.

#20

The leveling effect of water prevents differentiation of strong acids with pKa values below -1.74, as all strong acids are leveled to hydronium ions (H3O+).

Subject Specialist Commentary

Analytical perspective & practical exam advice from the Master10 academic board

Educator's Insight

Think of pKa as an intrinsic measure of how tightly an acid holds onto its hydrogen ions. While pH tells you how acidic a specific glass of liquid is at a given moment, pKa reveals the permanent chemical nature of the acid molecule itself. A smaller or negative pKa number means an acid gives up its protons readily, confirming it is a strong acid in aqueous solutions.

For competitive exams like UPSC and State PSC, never confuse solution pH with molecular pKa. Remember that at the half-equivalence point of a weak acid titration, pH equals pKa because acid and conjugate base concentrations match exactly. Examiners frequently test the carbonic acid blood buffer and the rule that electron-withdrawing halogens stabilize conjugate bases, which lowers pKa and raises overall acid strength.

What Is pKa? Acid Dissociation Constant & Henderson-Hasselbalch Equation

Key Concepts & Self-Assessment20 Key Facts

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