Ka / Kb / pKa Converter
Convert Ka, Kb, pKa, and pKb for conjugate acid-base pairs using Kw at 25 C.
What can you enter?
Choose the value you know. The calculator fills in the conjugate acid-base values and gives a strength hint.
| Known value | Meaning | Formula used |
|---|---|---|
| Ka | Acid dissociation constant | pKa = -log10(Ka) |
| Kb | Base dissociation constant | pKb = -log10(Kb) |
| pKa | Log form of Ka | Ka = 10^(-pKa) |
| pKb | Log form of Kb | Kb = 10^(-pKb) |
Formula used
pKb = -log10(Kb)
Ka * Kb = Kw = 1.0e-14 at 25 C
pKa + pKb = 14 at 25 C
Worked examples
Acetic acid: pKa = 4.76, so Ka = 10^-4.76 = 1.74e-5. The conjugate base has pKb = 14 - 4.76 = 9.24.
Ammonia: Kb = 1.8e-5, so pKb = -log10(1.8e-5) = 4.745. The conjugate acid has pKa = 9.255.
How to read acid and base strength
A larger Ka means a stronger acid. A smaller pKa also means a stronger acid because pKa is the negative logarithm of Ka. For bases, a larger Kb or smaller pKb means a stronger base.
| pKa range | Acid strength hint | What to remember |
|---|---|---|
| Below 0 | Very strong acid range | Often mostly dissociated in water |
| 0 to 2 | Strong acid range | Much stronger than typical weak acids |
| 2 to 7 | Weak acid range | Common for many carboxylic acids |
| 7 to 14 | Very weak acid range | Conjugate base is comparatively stronger |
Where this converter is useful
- Converting pKa table values into Ka for equilibrium calculations.
- Finding the conjugate base Kb from a weak acid Ka.
- Checking Henderson-Hasselbalch buffer problems.
- Comparing acid strength without switching between log and equilibrium formats by hand.
- Practicing Kw relationships at 25 C.
Common mistakes
- Thinking high pKa means stronger acid; it means weaker acid.
- Using natural log instead of base-10 log for pKa and pKb.
- Forgetting that Ka * Kb = Kw only applies to a conjugate acid-base pair.
- Using Kw = 1.0e-14 when the problem gives a different temperature.
- Rounding pKa too early before converting back to Ka.
Related Chemistry Tools
Ka / Kb / pKa Converter FAQs
When are Ka, Kb, pKa and pKb used?
These constants describe the strength of weak acids and bases. Ka is the acid dissociation constant (HA ■ H+ + A-), used for weak acids. Kb is the base dissociation constant (B + H2O ■ BH+ + OH-), used for weak bases. pKa = −log Ka and pKb = −log Kb convert these (often very small) numbers into more manageable figures. The smaller the pKa, the stronger the acid; smaller pKb means stronger base. Useful relation at 25 °C: pKa + pKb = 14 (for a conjugate acid–base pair). Ka · Kb = Kw = 10-14 ; pKa + pKb = 14
How to find Ka from pKa?
Just take the antilog: Ka = 10−pKa. Examples: acetic acid pKa = 4.76 → Ka = 10-4·76 ≈ 1.74 × 10-5; HF pKa = 3.17 → Ka = 10-3·17 ≈ 6.76 × 10-4. The smaller the pKa, the larger the Ka, and the stronger the acid. Knowing Ka lets you calculate pH of weak-acid solutions using the Ka expression and an ICE table. Ka = 10−pKa
Is low pKa a strong acid?
Yes — a low pKa indicates a strong acid. Lower pKa means higher Ka, which means the acid ionises to a greater extent in water. Strong acids like HCl have pKa ≈ −7 (extreme), nitric acid pKa ≈ −1.4, while weak acids like acetic acid have pKa ≈ 4.76 and very weak acids like phenol have pKa ≈ 9.95. As a rough rule, pKa < 0 indicates strong acids, pKa between 0 and 14 indicates weak acids.
How to calculate pKa from Ka?
Take the negative log: pKa = −log10Ka. Example: Ka for formic acid = 1.8 × 10-4 → pKa = −log(1.8 × 10-4) = 4 − log 1.8 = 4 − 0.255 = 3.745. So formic acid has pKa ≈ 3.75. The smaller this value, the stronger the acid. Always remember to use base-10 log, not natural log. pKa = − log10 Ka
How to find pKa from pH?
Use the Henderson–Hasselbalch equation rearranged: pKa = pH − log([A-]/[HA]). The cleanest way is at the half-equivalence point of a weak-acid–strong-base titration, where exactly half of HA has been neutralised, so [HA] = [A-], the log term is zero, and pKa = pH. So you can experimentally find pKa simply by reading off the pH at the half-equivalence point of the titration curve. This is one of the most elegant uses of the equation.
Does lower pKa mean more acidic?
Yes, exactly. Lower pKa = more acidic. Since pKa = −log Ka, a lower pKa implies a larger Ka, meaning the acid dissociates more in water and produces more H+ ions. Compare: HCl (pKa ≈ −7, completely ionised, very strong) > HF (pKa = 3.17, partial) > acetic acid (pKa = 4.76, weak) > water (pKa = 15.7, extremely weak).
What is pKa of water?
Water acts as both acid and base — it self-ionises. The auto-ionisation equilibrium is H2O ■ H+ + OH-, with Kw = 10-14 at 25 °C. However, when we strictly define Ka(H2O) accounting for the activity of water (55.5 M), we get Ka ≈ 1.8 × 10-16, giving pKa(H2O) ≈ 15.7. Some textbooks quote 14 by treating water concentration as part of Kw; the rigorous chemist's value is 15.7. Either way, water is a very weak acid.
Does high pKa mean stronger acid?
No, the opposite. A high pKa means a low Ka, which means the acid ionises only slightly in water — it is a weaker acid. So strong acids have low pKa (often negative), and weak acids have high pKa. Use this little ladder to keep things straight: HCl (pKa ≈ −7) is a strong acid; CH3COOH (pKa = 4.76) is weak; H2O (pKa = 15.7) is extremely weak; CH4 (pKa ≈ 50) is essentially not acidic at all.