Grade 11

Acid–base, redox, complexometric and precipitation titrations

Finding an unknown concentration by adding a solution of known concentration until the reaction is complete.

1 Lower secondary 2 Upper secondary 3 Undergraduate 4 Advanced 5 Research

IntuitionIntuition: add until the color changes

In a titration you add a solution of known concentration (the titrant) from a burette to a measured volume of the unknown solution until the reaction is exactly complete. That moment is the equivalence point; an indicator that changes color close to it (the end point) tells you when to stop.

25 mL of acid titrated with

\ce{NaOH}

of the same concentration, so the equivalence point is at 25 mL. Compare the strong acid

\ce{HCl}

with the weak acid

\ce{CH3COOH}

SchoolSchool level: the calculation

\ce{HCl + NaOH -> NaCl + H2O}

n(\ce{H+}) = n(\ce{OH-}) \;\Rightarrow\; C_a V_a = C_b V_b

Example: Titrating $\ce{HCl}$

20.0 mL of $\ce{HCl}$ of unknown concentration needs 25.0 mL of 0.100 mol/L $\ce{NaOH}$ to reach the equivalence point. Find $C_a$ .

Solution

$C_a = C_b V_b / V_a = 0.100 \times 25.0 / 20.0 = 0.125$ mol/L.

Drag the volume of

\ce{NaOH}

past 25 mL for the strong acid and watch phenolphthalein turn pink. Try other indicators and the weak acid; the corner shows the true pH.

UndergraduateUndergraduate: reading the curve

For a strong acid with a strong base the equivalence point is at $\text{pH} = 7$ and the jump is very steep. For the weak acid $\ce{CH3COOH}$ ( $K_a = 1.8\times10^{-5}$ ) the solution at equivalence contains acetate, a weak base, so the equivalence point lies above 7. At the half-equivalence point $\text{pH} = \text{p}K_a \approx 4.74$ , which is how $K_a$ is measured.

\text{pH} = \text{p}K_a + \log\frac{[\ce{A-}]}{[\ce{HA}]}