Henderson-Hasselbalch Calculator
Calculate buffer pH using the Henderson-Hasselbalch equation. Determine pH from pKa and concentration ratio, or solve for any variable. Includes buffer capacity and common buffer recipes.
pKa Calculator
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.
Common Acids
For % dissociation calculation
M
pKa⧉
4.760
Moderate weak acid
Ka⧉
1.738e-5
Acid dissociation constant
pKb (conjugate)⧉
9.240
Conjugate base strength
Kb (conjugate)⧉
5.754e-10
Base dissociation constant
% Dissociation⧉
1.31%
At 0.1 M concentration
Solution pH⧉
2.883
[H⁺] = 1.310e-3 M
Species Distribution vs. pH
pH 0.8
0% A⁻
pH 1.8
0% A⁻
pH 2.8
1% A⁻
pH 3.8
9% A⁻
pH 4.3
24% A⁻
pH 4.8
50% A⁻
pH 5.3
76% A⁻
pH 5.8
91% A⁻
pH 6.8
99% A⁻
pH 7.8
100% A⁻
pH 8.8
100% A⁻
HA (acid) A⁻ (conjugate base)
Acid Strength Reference
| Acid | Formula | pKa₁ | pKa₂ | pKa₃ | Type | Relative Strength |
|---|---|---|---|---|---|---|
| Hydrochloric acid | HCl | -7 | — | — | Strong | |
| Sulfuric acid | H₂SO₄ | -3 | 1.99 | — | Strong/Moderate | |
| Nitric acid | HNO₃ | -1.4 | — | — | Strong | |
| Phosphoric acid | H₃PO₄ | 2.15 | 7.2 | 12.35 | Triprotic | |
| Hydrofluoric acid | HF | 3.17 | — | — | Weak | |
| Formic acid | HCOOH | 3.75 | — | — | Weak | |
| Ascorbic acid (Vit C) | C₆H₈O₆ | 4.1 | 11.6 | — | Diprotic | |
| Benzoic acid | C₆H₅COOH | 4.2 | — | — | Weak | |
| Acetic acid | CH₃COOH | 4.76 | — | — | Weak | |
| Carbonic acid | H₂CO₃ | 6.35 | 10.33 | — | Diprotic | |
| Hydrogen sulfide | H₂S | 7 | 14 | — | Diprotic | |
| Hypochlorous acid | HOCl | 7.53 | — | — | Weak | |
| Boric acid | B(OH)₃ | 9.24 | — | — | Weak | |
| Ammonium ion | NH₄⁺ | 9.25 | — | — | Conjugate | |
| Hydrocyanic acid | HCN | 9.31 | — | — | Weak | |
| Phenol | C₆H₅OH | 10 | — | — | Weak | |
| Water | H₂O | 15.7 | — | — | Amphoteric | |
| Ethanol | C₂H₅OH | 16 | — | — | Very weak | |
| Acetone | CH₃COCH₃ | 19.3 | — | — | Very weak |
Planning notes, formulas, and examples
About the pKa Calculator
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.
When This Page Helps
Convert between pKa and Ka, find conjugate base properties, calculate dissociation at any concentration, and compare acid strengths with a comprehensive reference table.
How to Use the Inputs
- Enter a pKa value (or Ka) to convert between them.
- View the conjugate base pKb and Kb automatically.
- Enter an acid concentration to see percent dissociation.
- Select from preset acids to load their pKa values.
- Compare multiple acids in the ranked strength table.
- Toggle polyprotic acid mode to see all ionization steps.
Formula used
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 = 0Example Calculation
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.
Tips & Best Practices
- At pH = pKa, the acid is exactly 50% dissociated.
- The conjugate base is always weaker: pKb = 14 - pKa (at 25°C).
- For the Henderson-Hasselbalch equation, you need pKa. For bases, use pKa of the conjugate acid.
- Polyprotic acids: Ka values decrease by ~10⁵ for each successive ionization.
- Strong acids (pKa < -1) are "leveled" in water — all appear equally strong.
- In organic chemistry, pKa guides deprotonation selectivity: always deprotonate the most acidic H first.
pKa in Organic Chemistry
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).
pKa and Drug Absorption
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.
Measuring pKa
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.
Sources & Methodology
Last updated:
Frequently Asked Questions
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.
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