Article — pKa Calculator
pKa Calculator: Ka, pKa, and Buffer pH
pKa is the negative base-10 logarithm of the acid dissociation constant Ka. A pKa of 4.76 means Ka = 1.8 × 10−5, which describes acetic acid. Lower pKa values mean stronger acids; HCl sits around −7 while phenol sits at 10.
Søren Sørensen introduced the logarithmic acidity scale at the Carlsberg Laboratory in 1909. The pKa extension followed soon after and gave chemists a single number that captures how readily a weak acid donates its proton in water at 25 °C.
What is pKa?
pKa quantifies acid strength on a manageable logarithmic scale. The underlying equilibrium is HA ⇌ H+ + A−, with the equilibrium constant Ka = [H+][A−]/[HA]. Ka values span 20 orders of magnitude for ordinary acids, which is why the logarithmic form is so useful.
A pKa below 0 means the acid is essentially fully dissociated in dilute water (a strong acid). Between 0 and 3, the acid is moderate. From 3 to 10 it is weak, and above 10 it is very weak. Each unit of pKa is a factor of 10 in Ka, so a pKa difference of 3 means a 1000-fold difference in dissociation.
The "p" in pKa, pH, and pKw all come from the German Potenz, meaning "power" or "exponent". Sørensen picked the letter in 1909 to mean "the negative logarithm of."
pKa vs pH: what is the difference?
pH is a property of a solution; pKa is a property of an acid. The pH of vinegar depends on how much acetic acid you dissolved, but the pKa of acetic acid is always 4.76 at 25 °C, whether you have 1 mL or 1 liter. Confusion between the two is one of the most common errors in introductory chemistry.
The two are linked: at the half-equivalence point of a titration, pH = pKa exactly. That is also where a buffer made from equal parts weak acid and conjugate base has its maximum buffering capacity.
The pKa formula and its inverse
Two equations cover most pKa work:
pKa = −log₁₀(Ka) convert Ka → pKaKa = 10^(−pKa) convert pKa → KapH = pKa + log₁₀([A⁻]/[HA]) Henderson-HasselbalchpKa + pKb = 14 conjugate pair at 25 °CFor acetic acid: pKa = −log10(1.8 × 10−5) = 4.74. The calculator uses the same one-line operation, plus the inverse, plus the Henderson-Hasselbalch form for buffer calculations.
pKa values of common acids
The table built into the calculator covers 12 acids spanning the strong-to-very-weak range. A few worth memorizing:
- Acetic acid (vinegar) = 4.76 — defines the classic acetate buffer
- Phosphoric acid = 2.14, 7.20, 12.37 — three pKa values for the three protons
- Carbonic acid = 6.35 — central to blood pH regulation
- Ammonium ion = 9.25 — gives the NH3/NH4+ buffer at pH 9 to 10
- Hydrofluoric acid = 3.17 — weak despite being highly reactive
- Phenol = 10.0 — weakly acidic, useful in organic synthesis
- Hydrochloric acid = −7 — fully dissociated in dilute water
Henderson-Hasselbalch and buffer pKa
The Henderson-Hasselbalch equation predicts buffer pH from pKa and the ratio of conjugate base to weak acid: pH = pKa + log10([A−]/[HA]). When the ratio is 1, the logarithm is 0 and pH equals pKa. A 10-to-1 ratio shifts pH by exactly one unit above or below pKa.
A buffer is effective only within pKa ± 1. Past that window, the [A−]/[HA] ratio becomes too extreme and the buffer cannot absorb additional acid or base. This is why blood relies on multiple buffers (bicarbonate, phosphate, protein) — no single pKa covers all conditions.
The Henderson-Hasselbalch equation assumes the dissociation is small relative to total concentration. At very low concentrations (below 1 mM) or for moderately strong acids (pKa under 2), the approximation breaks down and you need the full ICE-table solution to the quadratic equation.
pKa + pKb = 14 and the conjugate base
For any conjugate acid-base pair in water at 25 °C, pKa + pKb = pKw = 14. Acetic acid (pKa 4.76) has a conjugate base — the acetate ion — with pKb = 9.24. A strong acid produces a very weak conjugate base, which is why the chloride ion is effectively non-basic in water.
At 37 °C (body temperature), pKw drops to about 13.6, so the neutral pH of pure water shifts to 6.8. This matters in biochemistry: enzymatic optima reported at room temperature need adjustment for in vivo conditions.
pKa in drug ionization and absorption
Drug pKa governs how much of a dose exists in the un-ionized, lipid-soluble form at each gastrointestinal pH. Aspirin (pKa 3.5) is mostly un-ionized in the stomach (pH 2) and absorbs rapidly there. Morphine (pKa 8) is mostly ionized in the stomach and absorbs in the small intestine where pH rises to 6 or 7.
The Henderson-Hasselbalch equation predicts the fraction ionized: at pH = pKa, half the drug is ionized. Each unit difference between pH and pKa shifts the ratio by a factor of 10. Formulators use this to design drugs that hit specific absorption windows.
For a weak acid drug, ionization rises as pH rises above pKa. For a weak base drug, ionization rises as pH drops below pKa. The un-ionized form is the absorbable form across most membranes.
Common pKa mistakes to avoid
The most frequent error is reversing the direction: students assume that a higher pKa means a stronger acid because it sounds like pH. The opposite is true — higher pKa means a weaker acid. The second common error is ignoring temperature. Tabulated values are at 25 °C; at 37 °C every pKa shifts slightly, and so does pKw.
A third trap is forgetting that polyprotic acids have multiple pKa values. Phosphoric acid has three (2.14, 7.20, 12.37) and the dominant species depends on pH. At pH 7, the buffer system is mostly H2PO4−/HPO42−, and only pKa2 matters.
- Reversed scale — low pKa = strong acid, not the other way around
- Forgetting temperature — pKa values shift with T; 25 °C is the convention
- Buffer outside range — Henderson-Hasselbalch fails past pKa ± 1
- Ignoring polyprotic pKa2, pKa3 — phosphate buffer uses pKa2, not pKa1
- Confusing pKa with pKw — pKw = 14 is the water equilibrium, separate from acid pKa