Describe how to identify the end-point of a titration using an indicator

Acid‑Base Titrations – Identifying the End‑Point with an Indicator

Learning Objectives (AO1, AO2, AO3)

• Describe the purpose of each piece of apparatus used in a titration.
• Explain how an indicator allows the end‑point to be detected and why this approximates the equivalence point.
• Choose a suitable indicator for a given acid‑base system.
• Carry out a titration safely, record observations correctly and evaluate sources of error.
• Calculate the unknown concentration of an analyte from the titration data.

1. What Is a Titration?

A titration is a quantitative analytical technique in which a solution of known concentration (the titrant) is added gradually to a measured volume of another solution (the analyte) until the reaction is complete. The volume of titrant required is then used to calculate the unknown concentration of the analyte.

The point at which the stoichiometric amounts of acid and base have reacted is the equivalence point. In practice we detect the nearby end‑point by a colour change of an indicator. Because the indicator changes colour over a narrow pH range that overlaps the pH at the equivalence point, the end‑point is a good approximation of the equivalence point (e.g. for a strong acid / strong base the equivalence pH ≈ 7, which lies within the transition range of phenolphthalein).

2. Essential Apparatus and Their Purpose

3. Choosing the Right Indicator

The indicator’s pH transition range must overlap the pH that the solution will have at the equivalence point. Use the decision‑tree below to select an indicator for the most common acid‑base systems.

4. General Procedure (AO3 – Experimental Skills)

1. Preparation
   • Wear safety goggles, lab coat and gloves.
   • Rinse the burette with a small amount of the titrant, then fill it to the zero mark. Record the initial volume (to two decimal places).
   • Using a calibrated volumetric pipette, transfer the required volume of the analyte into a clean conical flask. If the sample is viscous, add a few drops of distilled water to aid transfer.
   • Add 2–3 drops of the chosen indicator. Swirl gently to mix uniformly.
2. Titration
   • Place the flask on a white tile.
   • Open the burette tap and add titrant while constantly swirling the flask.
   • When you are within about 1 cm³ of the expected end‑point, add the titrant drop‑wise (≈0.05 cm³ per drop).
   • Observe the colour change. The first permanent colour that persists for at least 30 s (even if a couple of extra drops are added) is taken as the end‑point. Record the final burette reading (to two decimal places).
3. Recording Data (AO2)
   • Initial burette volume, final burette volume, volume of titrant used.
   • Temperature of the room (≈ 25 °C) – note any deviation.
   • Indicator used and number of drops.
   • Any observations of colour intensity or difficulty in detecting the change.

5. Recognising a True End‑Point

• Sharp, distinct colour change – the new colour should appear suddenly, not gradually.
• Persistence – the colour must remain unchanged for at least 30 s.
• Reproducibility – repeating the titration should give the same end‑point volume within experimental error (±0.05 cm³).

6. Common Sources of Error & How to Minimise Them

7. Units, Significant Figures and Reporting Results (AO2)

• Volumes from the burette are recorded to two decimal places (e.g., 23.47 cm³).
• Concentrations are reported to the appropriate number of significant figures, usually three for IGCSE work (e.g., 0.102 M).
• When performing calculations, keep extra decimal places internally and round only in the final answer.

8. Example Calculations

Example 1 – 1 : 1 Stoichiometry (Strong Acid / Strong Base)

Problem: 25.00 cm³ of an unknown HCl solution is titrated with 0.100 M NaOH using phenolphthalein. The burette reading changes from 5.00 cm³ to 30.25 cm³ at the end‑point.

1. Volume of NaOH used:
   \(V_{\text{NaOH}} = 30.25 - 5.00 = 25.25\ \text{cm}^3\)
2. Moles of NaOH added:
   \(n_{\text{NaOH}} = 0.100\ \text{mol L}^{-1} \times 25.25 \times 10^{-3}\ \text{L}=2.525 \times 10^{-3}\ \text{mol}\)
3. Reaction stoichiometry (1 mol HCl : 1 mol NaOH):
   \(n_{\text{HCl}} = n_{\text{NaOH}} = 2.525 \times 10^{-3}\ \text{mol}\)
4. Concentration of HCl:
   \(C_{\text{HCl}} = \dfrac{2.525 \times 10^{-3}\ \text{mol}}{25.00 \times 10^{-3}\ \text{L}} = 0.101\ \text{M}\) (3 s.f.)

Example 2 – 1 : 2 Stoichiometry (Strong Base / Diprotic Acid)

Problem: 20.00 cm³ of 0.050 M H₂SO₄ is titrated with 0.100 M NaOH using phenolphthalein. The end‑point occurs at a final burette reading of 35.20 cm³ (initial reading 2.00 cm³).

1. Volume of NaOH used:
   \(V_{\text{NaOH}} = 35.20 - 2.00 = 33.20\ \text{cm}^3\)
2. Moles of NaOH added:
   \(n_{\text{NaOH}} = 0.100\ \text{mol L}^{-1} \times 33.20 \times 10^{-3}\ \text{L}=3.32 \times 10^{-3}\ \text{mol}\)
3. Reaction stoichiometry: 1 mol H₂SO₄ reacts with 2 mol NaOH.
   \(n_{\text{H₂SO₄}} = \dfrac{n_{\text{NaOH}}}{2}= \dfrac{3.32 \times 10^{-3}}{2}=1.66 \times 10^{-3}\ \text{mol}\)
4. Check the given concentration:
   Expected moles of H₂SO₄ = 0.050 M × 20.00 cm³ = 1.00 × 10⁻³ mol.
   The larger calculated value shows that the end‑point was overshot; this illustrates the importance of careful colour observation.

9. Safety & Record‑Keeping Checklist (AO3)

• Wear goggles, lab coat and gloves.
• Rinse burette with titrant before filling.
• Record initial and final burette readings to two decimal places.
• Note temperature of the room (≈ 25 °C).
• Use a white tile to aid colour detection.
• Dispose of waste according to school regulations.

10. Summary

• The end‑point is detected by a colour change of an indicator whose pH transition overlaps the equivalence‑point pH.
• Accurate measurement of the analyte (pipette) and titrant (burette) is essential for reliable results.
• Select the indicator using the decision‑tree or the table of typical pH ranges.
• Observe the colour change carefully, ensuring it is sharp, permanent and reproducible.
• Identify and minimise common sources of error (parallax, over‑titration, temperature, excess indicator).
• Record all data with the correct number of significant figures and use the stoichiometry of the reaction to calculate the unknown concentration.