# Why there is an increase in the pH at the starting of a titration between a weak acid and a strong base?

Why there is increase in the pH at the starting of a titration between a weak acid and a strong base even though there is acid already present to neutralize it? My intuition says that there should not be even a slight increase in pH until all of the weak acid is dissociated and neutralized by the base

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• What does the equation pH = pKa + log([A-]/[HA]) say to you? Jan 22 at 12:13
• @Poutnik . more acid will dissociate as pH increases. Jan 22 at 12:35
• @NikhilVerma It has also the other side - increasing acid dissociation ratio by acid neutralization increases pH, Jan 22 at 14:55

Let's start from the basic formula : $$\mathrm{p}H = \mathrm{p}K_a + \log\mathrm{\frac{[A^-]}{[HA]} = \mathrm{p} K_a + {log{\frac{\mathrm{[A_o]} + n}{[HA]_o - n}}}}$$ The derivate of log$$x$$ is $$1/x$$. So when introducing $$n$$ mole $$\ce{NaOH}$$, we get a curve with a slope equal to : $$\frac{d\mathrm pH}{dn} = \mathrm{\frac{[HA]}{[A^-]}} = \frac{[\mathrm{H^+]}}{K_{a}}$$
At the beginning of the titration, when $$n$$ is small, the acidic concentration $$\ce{[H^+]}$$ is high. So the slope of the curve is steep. When $$n$$ grows bigger, $$\ce{[H+]}$$ becomes smaller. As a consequence, the slope of the curve decreases.