Chemistry Labs
Upper secondary · 15 min

Acid–base titration with an indicator

Drip NaOH from a burette into an acid and watch the indicator change color at the equivalence point.

3D burette above an Erlenmeyer flask: the liquid in the flask changes color as NaOH is added, with pH and volume shown in the corner.

Goal

Find the volume of NaOH\ce{NaOH} at the equivalence point, see how the indicator responds, and compare a strong with a weak acid.

Apparatus and reagents

25 mL of acid (0.10 mol/L) in an Erlenmeyer flask; 0.10 mol/L NaOH\ce{NaOH} in a 50 mL burette; a few drops of indicator.

Procedure

  1. Start with the strong acid HCl\ce{HCl} and phenolphthalein. Move the volume slider slowly from 0 mL and watch the color and the pH.
  2. Go slowly between 24 and 26 mL and note the volume at which the color first persists.
  3. Switch the acid to the weak CHX3COOH\ce{CH3COOH} and compare the pH at 25 mL.
  4. Keep the weak acid and change the indicator to methyl orange: at what volume does it change color?

What to observe

  • For HCl\ce{HCl} the pH starts at 1.00, is 3.70 at 24.9 mL, 7.00 at 25.0 mL and 10.30 at 25.1 mL: a jump of about 6.6 units from one drop of NaOH\ce{NaOH}.
  • For CHX3COOH\ce{CH3COOH} the pH starts near 2.88 and at 25.0 mL is about 8.72, not 7, because acetate is a weak base.
  • Phenolphthalein changes near the equivalence point for both acids; methyl orange changes far too early with the weak acid.

Explanation

At the equivalence point n(HX+)=n(OHX−)n(\ce{H+}) = n(\ce{OH-}), so CaVa=CbVbC_aV_a = C_bV_b: with equal concentrations the equivalence is at 25.0 mL. An indicator is useful when its color-change range lies on the steep part of the curve near the equivalence pH. Methyl orange (pH 3.1–4.4) is passed at about 12 mL with the weak acid, long before the equivalence point at 25 mL, so it would give a badly wrong result.

History of the experiment

Volumetric analysis began with Gay-Lussac's 1829 titration of bleach with indigo, and by mid-century Descroizilles' burette and Mohr's pipette had made 'titration' a routine industry tool. Arrhenius's ionic theory (1887) and Ostwald's indicator chemistry explained why the colour snaps exactly when n(HX+)=n(OHX−)n(\ce{H+}) = n(\ce{OH-}) — the equivalence this lab hunts drop by drop.

Chemists behind it

Svante ArrheniusWilhelm OstwaldJoseph Louis Gay-Lussac

Related topics

Acid–base, redox, complexometric and precipitation titrationsElectrolytes, pH, buffer solutions

Virtual experiment: a simplified model to build intuition. It does not replace real lab work or safety training; never repeat chemistry at home without supervision.