Convert between pKa and Ka, calculate conjugate pKb, predict acid strength, and compare acid-base properties. Includes a database of common acids with pKa values and polyprotic acid support.
The pKa of an acid is the negative logarithm of its acid dissociation constant Ka: pKa = -log₁₀(Ka). It quantifies acid strength — the lower the pKa, the stronger the acid. This single number is perhaps the most useful quantity in acid-base chemistry, affecting drug design, buffer preparation, titration endpoints, enzyme activity, and organic reaction mechanisms.
Every Brønsted acid HA has a conjugate base A⁻ with its own basicity constant Kb. The two are linked: pKa + pKb = pKw = 14.00 at 25°C. A strong acid (low pKa) has a weak conjugate base (high pKb), and vice versa. This reciprocal relationship is fundamental to understanding acid-base equilibria.
This calculator converts between pKa and Ka, calculates the conjugate pKb and Kb, estimates percent dissociation at a given concentration, predicts whether a species acts as acid or base in water, and provides a searchable reference table of over 30 common acids ranked by strength. Polyprotic acids show each ionization step.
Convert between pKa and Ka, find conjugate base properties, calculate dissociation at any concentration, and compare acid strengths with a comprehensive reference table. This pka calculator helps you compare outcomes quickly and reduce avoidable mistakes when making day-to-day care decisions. Use the estimate as a planning baseline and confirm final decisions with a qualified professional when risk is high.
pKa = -log₁₀(Ka) Ka = 10^(-pKa) pKa + pKb = pKw = 14.00 (at 25°C) Ka × Kb = Kw = 1.0 × 10⁻¹⁴ % Dissociation = (α × 100) where α = [H⁺]/C₀ from quadratic: x² + Ka·x - Ka·C = 0
Result: Ka = 1.74 × 10⁻⁵, pKb = 9.24, 1.33% dissociated
Acetic acid (pKa = 4.76): Ka = 10⁻⁴·⁷⁶ = 1.74 × 10⁻⁵. Conjugate base: pKb = 14.00 - 4.76 = 9.24. At 0.1 M: x = √(Ka × C) ≈ 1.32 × 10⁻³, so 1.32% dissociated with pH = 2.88.
Acidity trends follow electronegativity, hybridization, resonance, and inductive effects. sp hybridized C-H bonds (pKa ~25) are more acidic than sp² (~44) or sp³ (~50). Electronegative substituents lower pKa through inductive withdrawal: trifluoroacetic acid (0.23) vs acetic acid (4.76). Resonance stabilization of the conjugate base dramatically lowers pKa: phenol (10.0) vs cyclohexanol (18).
The pH partition hypothesis states that only un-ionized drug molecules cross biological membranes efficiently. Stomach acid (pH 1-2) favors absorption of weak acids (e.g., aspirin, pKa 3.5), while the intestine (pH 6-8) better absorbs weak bases (e.g., morphine, pKa 8.0). Knowing the drug's pKa lets you predict absorption site and bioavailability.
pKa is measured by potentiometric titration, spectrophotometric methods, or NMR titration. Potentiometric titration monitors pH during addition of strong base, with pH = pKa at the half-equivalence point. Spectrophotometric methods track UV-vis absorbance changes as the acid-base ratio shifts, useful for sparingly soluble compounds.
A negative pKa means Ka > 1, indicating a strong acid that dissociates essentially completely. HCl has pKa ≈ -7, meaning Ka ≈ 10⁷. Most strong acids have pKa < -1.
Lower pKa = stronger acid. Each unit decrease in pKa means the acid is 10× stronger (Ka is 10× larger). Sulfuric acid (pKa ≈ -3) is ~10⁸ times stronger than acetic acid (pKa = 4.76).
Yes. Extremely weak acids like methane (pKa ~ 48) or ethane (pKa ~ 50) have very high pKa values. These can only be deprotonated by extraordinarily strong bases in aprotic solvents.
Polyprotic acids lose protons sequentially. Ka₁ > Ka₂ always (the second proton is harder to remove from a negatively charged species). For H₂SO₄: Ka₁ is very large (strong), Ka₂ = 0.012.
Most carboxylic acid pKa values increase slightly with temperature. Amine pKa values decrease significantly (e.g., Tris drops ~0.03 per °C). Phosphate buffers are relatively temperature-stable.
A drug's pKa determines its ionization at physiological pH, affecting membrane permeability, solubility, and bioavailability. The Henderson-Hasselbalch equation predicts the fraction ionized at any pH.