Properties of Buffers

A buffer holds two reservoirs at once: a weak acid (HA) and its conjugate base (A-), both present in significant amounts. Add strong acid, and the A- reservoir soaks up the extra H+, turning into HA. Add strong base, and the HA reservoir gives up a proton to neutralize it, turning into A-. Either way, the ratio of HA to A- shifts only a little, and since pH depends on the log of that ratio (Henderson-Hasselbalch: pH = pKa + log([A-]/[HA])), a small ratio shift means a small pH shift. Without both species present, there's nothing to catch the intruder.

Picture a flask with, say, 0.40 mol acetic acid and 0.40 mol acetate. Squirt in 0.05 mol NaOH: it reacts completely with 0.05 mol of the acetic acid, leaving 0.35 mol acetic acid and 0.45 mol acetate. The ratio barely moved from 1:1 to 0.78:1, so the pH barely moved from pKa. That's the whole mechanism: strong acid converts conjugate base into weak acid mole-for-mole, strong base does the reverse, and you plug the new amounts back into Henderson-Hasselbalch.

This is why HCl and NaCl never make a buffer: HCl ionizes completely, so there's no leftover molecular HA to react with added base, and Cl- is such a weak base it won't touch added acid. Blood uses the carbonic acid/bicarbonate pair (H_2CO₃ / HCO₃-) the same way: CO₂ from exercising muscle forms H_2CO₃, but bicarbonate absorbs the extra protons fast enough to hold pH near 7.4, linking this topic directly to the titration curves and equivalence points covered elsewhere in the unit.