A buffer pairs a weak acid HA with its conjugate base A⁻. The Henderson-Hasselbalch equation, pH = pKa + log₁₀([A⁻]/[HA]), says the pH depends only on the pKa and the ratio of the two forms — not on how concentrated the buffer is. Move the ratio slider and the beaker's pH follows this formula exactly.
Adding strong base converts HA into A⁻ one mole at a time (HA + OH⁻ → A⁻ + H₂O). While both species remain present in reasonable amounts, this barely moves the pH — the buffering region, roughly pKa ± 1. Once essentially all HA is consumed (100% on the titrant slider), there is nothing left to soak up further base and the pH jumps sharply: the equivalence point.
Buffer capacity β = 2.303·C·(Ka·[H⁺])/(Ka+[H⁺])² measures how many moles of strong acid or base per liter it takes to shift the pH by one unit. It peaks exactly at pH = pKa (where HA and A⁻ are equal) and collapses toward the equivalence point — which is why buffers are chosen so their pKa sits near the pH they need to defend, the same principle blood uses (pKa 6.1 bicarbonate buffer) to hold pH near 7.4.