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

What is the End‑Point?

The end‑point is the moment in a titration when the amount of titrant added just equals the amount of analyte in the solution. Think of it like filling a cup to the brim – the moment the liquid starts to spill is the end‑point. In a titration, we want to stop just before the solution “spills” into the next pH range.

Why Use an Indicator?

An indicator is a weak acid or base that changes colour at a particular pH range. It acts like a traffic light: green means the reaction is still moving, red means it’s finished. The colour change tells us the end‑point without needing a pH meter.

Choosing the Right Indicator

Pick an indicator whose colour change pH range straddles the expected equivalence point of the reaction. For a strong acid–strong base titration, the equivalence point is at pH 7, so phenolphthalein (pH 8.2–10) is a good choice. For weak acid–strong base, use an indicator that changes near the higher pH of the equivalence point.

Step‑by‑Step: Detecting the End‑Point

1. Fill the burette with the titrant (e.g., NaOH) and record the initial volume.
2. Place the analyte (e.g., HCl) in a clean beaker and add 2–3 drops of the chosen indicator.
3. Start the titration, adding the titrant slowly while swirling the beaker.
4. Watch the colour change. When the colour shifts and stays for a few seconds, the end‑point is reached.
5. Stop the titration, record the final burette reading, and calculate the volume used.

⚠️ Tip: Add the titrant in small increments near the expected end‑point to avoid overshooting.

Exam Tip Box

• Remember the pH range: Match the indicator’s colour change to the expected equivalence pH. • Describe the colour change: In your answer, state the colour before and after the end‑point (e.g., “colorless to pink”). • Explain why it works: Mention that the indicator’s weak acid/base properties cause a rapid shift in colour at a specific pH. • Use diagrams: A simple pH‑indicator curve can illustrate the concept. • Practice calculations: Be comfortable converting volume to moles and vice versa.

Quick Maths Check

If you added 25.00 mL of 0.100 M NaOH to 20.00 mL of 0.100 M HCl, how many moles of NaOH were used?

$$n_{\text{NaOH}} = C \times V = 0.100\,\text{mol L}^{-1} \times 0.02500\,\text{L} = 0.00250\,\text{mol}$$

Since the reaction is 1:1, the same number of moles of HCl were neutralised.

Analogy Corner

Think of the titration as a game of “balance the scales”. The indicator is the balancing scale’s needle. When the needle stops moving (colour change), the scales are balanced – that’s your end‑point! 🎯