Strong Acid-Strong Base Titration Curves

Start with 25 mL of 0.10 M HCl in a flask and add 0.10 M NaOH drop by drop. Early on, you have a huge excess of H+ sitting in solution. Each small addition of base neutralizes a tiny fraction of that excess, so the concentration of leftover H+ barely changes, and neither does the pH. But as you approach the equivalence point, the amount of H+ left is shrinking toward zero, and the same size drop of base now wipes out a much larger fraction of what remains. A concentration that drops from 0.001 M to 0.0000001 M only looks like a small change in moles, but pH = -log[H+] turns that into a jump of four whole units. That log compresses huge changes at low concentration into small numbers, which is exactly why the curve looks flat then nearly vertical.

To find pH before equivalence, subtract moles of base added from original moles of acid to get leftover moles of H+, divide by total volume, then take -log. At exactly 25.0 mL (equivalence for 25 mL of 0.10 M acid titrated with 0.10 M base), moles of acid and base are equal, all the H+ has been consumed, and the solution is just water and a neutral salt, giving pH 7. Past that point, excess OH- builds up and you solve for pOH instead, then convert.

This curve sets up everything after it: weak acid-strong base titrations start higher, have a buffer region, and reach equivalence above pH 7, all because the acid is not fully dissociated to begin with.