Acid-Base & pH Calculator

An acid is a proton (H⁺) donor and a base is a proton acceptor — the Brønsted–Lowry definition, and the one every panel on this page uses. Vinegar (acetic acid) donates a proton to water; ammonia accepts one and becomes ammonium. Every acid is paired with the species left behind after it gives up its proton: that pair — acetic acid CH₃COOH with acetate CH₃COO⁻, or ammonium NH₄⁺ with ammonia NH₃ — is a conjugate acid–base pair, and the whole of buffer and titration chemistry is bookkeeping of those pairs.

Because [H⁺] spans more than fourteen powers of ten between battery acid and drain cleaner, chemists compress it: pH = −log₁₀[H⁺]. Water itself always splits a little, [H⁺][OH⁻] = K_w = 1.0×10⁻¹⁴ at 25°C, which is why pH + pOH = 14: raise one and the other falls by the same amount. Neutral simply means [H⁺] = [OH⁻] = 10⁻⁷, which is pH 7 only at 25°C — hot pure water sits slightly below 7 and is still perfectly neutral. As a scale check the page can reproduce: pH 3.5 (about the acidity of orange juice) corresponds to [H⁺] = 3.16×10⁻⁴ mol/L.

Strong acids such as HCl dissociate essentially completely — pour them in water and virtually every molecule hands over its proton. Weak acids reach an equilibrium: HA ⇌ H⁺ + A⁻, described by K_a = [H⁺][A⁻]/[HA]. Acetic acid has K_a = 1.8×10⁻⁵, usually quoted as pK_a = 4.74 (the same −log trick as pH). The useful part is the pairing rule: K_a × K_b = K_w for a conjugate pair, so a stronger acid automatically has a weaker conjugate base. Acetate, the partner of weak acetic acid, is a very weak base with K_b = 5.6×10⁻¹⁰ (pK_b = 9.26) — both conversions run in the Ka & Kb panel above.

A buffer is a weak acid and its conjugate base present in comparable amounts. Added H⁺ gets eaten by A⁻; added OH⁻ gets eaten by HA; either way the ratio barely moves, so the pH barely moves. The bookkeeping collapses into pH = pK_a + log₁₀([A⁻]/[HA]). When the two concentrations are equal, pH = pK_a exactly, and every ten-fold shift in the ratio moves pH by exactly one unit — the 1:10 sample button gives pH 5.76 for an acetate buffer whose pK_a is 4.76. Reliable buffering only spans pK_a ± 1 (ratio 0.1–10), which is why choosing a buffer means choosing one whose pK_a sits near your target pH. Blood is the showcase: the bicarbonate/CO₂ pair holds plasma at 7.35–7.45, and enzymes stop working outside that window.

Adding strong base to an acid changes the [A⁻]/[HA] ratio continuously, and the curve is the Henderson–Hasselbalch equation drawn out. Early on the buffer soaks up the base and the pH crawls — the flat region. At the equivalence point the base exactly matches the acid moles; there the curve is nearly vertical, which is why that point flags the end of a titration. Halfway there, [A⁻] = [HA], so the pH equals the pK_a — a standard way to measure pK_a from a curve. Strong acid + strong base meets at pH 7, but weak acid + strong base finishes above 7 (near 8.7 for the default acetic-acid sample), because what remains in the flask is a weak base.

• Strong means concentrated. No — strength is how completely an acid dissociates, concentration is how much of it there is. 0.001 M HCl is a dilute strong acid; glacial acetic acid is a concentrated weak one.
• Neutral is always pH 7. Only at 25°C, where K_w = 10⁻¹⁴. Neutral means [H⁺] = [OH⁻], full stop.
• pH stops at 0 and 14. Those are just the boundaries of the 25°C water scale. 12 M HCl is around pH −1, and concentrated NaOH exceeds 14.
• Weak acid pH = −log of its concentration. That shortcut works only for strong acids. A weak acid's pH also needs its K_a, because most molecules never dissociate.

Related tools: Chemical Equilibrium (where equilibrium constants like K_a come from), Spectroscopy (measuring concentrations with light), and Bio Calculators (buffer recipes in micromolar units).