Titration Curves
1. Sketch the following graphs; state and explain whether each graph can be used in a titration; and state an appropriate indicator for each titration from the table below.
| Indicator | pH range |
|---|---|
| Bromocresol green | 3.8 – 5.4 |
| Bromothymol blue | 6.0 – 7.6 |
| Thymol blue | 8.0 – 9.6 |
Suitable Indicator: All three will work.
This is a Strong Acid / Strong Base titration. The vertical section of the curve typically spans from pH 3 to pH 10. Since the pH ranges of all three indicators (Bromocresol green, Bromothymol blue, and Thymol blue) fall within this vertical section, any of them would change colour sharply at the equivalence point.
Suitable Indicator: Thymol Blue
Explanation: This is a titration of a Weak Acid with a Strong Base. The salt formed at the equivalence point is Sodium Ethanoate, which is basic. Consequently, the equivalence point will be at a pH greater than 7 (typically pH 8-9). The indicator must change colour during the vertical section of the curve, which occurs in this alkaline region. Thymol Blue (pH 8.0 – 9.6) matches this range. Bromocresol green is incorrect as it changes colour in the acidic region (pH 3.8 – 5.4), where the solution is still buffering.
The pH change is too gradual at the equivalence point so colour change of indicator is difficult to judge because they are a weak acid and a weak base.
In Weak Acid / Weak Base titrations, there is no distinct vertical section on the pH curve. This means the pH changes slowly throughout the entire addition, making it impossible for an indicator to show a sharp end-point.
2. 0.5mol dm⁻³ Sodium hydroxide was added dropwise to a conical flask of a weak acid. The pH was determined using a pH meter. The results are in the table below.
| Volume OH⁻ / cm³ | pH |
|---|---|
| 0.00 | 2.49 |
| 2.00 | 3.61 |
| 4.00 | 3.95 |
| 6.00 | 4.16 |
| 8.00 | 4.32 |
| 10.00 | 4.46 |
| 12.00 | 4.58 |
| 14.00 | 4.7 |
| 16.00 | 4.82 |
| 18.00 | 4.94 |
| 20.00 | 5.06 |
| 22.00 | 5.2 |
| 24.00 | 5.36 |
| 26.00 | 5.57 |
| 28.00 | 5.91 |
| Volume OH⁻ / cm³ | pH |
|---|---|
| 30.00 | 9.07 |
| 32.00 | 12.24 |
| 34.00 | 12.53 |
| 36.00 | 12.69 |
| 38.00 | 12.8 |
| 40.00 | 12.89 |
| 42.00 | 12.95 |
| 44.00 | 13.01 |
| 46.00 | 13.05 |
| 48.00 | 13.09 |
| 50.00 | 13.12 |
| 52.00 | 13.15 |
| 54.00 | 13.18 |
| 56.00 | 13.21 |
| 58.00 | 13.23 |
Phenolphthalein.
This is because phenolphthalein has a pH range of 9-10, which is inside the vertical section of the pH curve shown by the data.
Looking at the data, the largest pH jump occurs between 28 cm³ (pH 5.91) and 32 cm³ (pH 12.24). The center of this vertical section is at 30.00 cm³ (pH 9.07).
Half-neutralisation occurs at exactly half the equivalence volume.
Volume at ½ neutralisation = 30 cm³ / 2 = 15.00 cm³.
From the table:
At 14.00 cm³, pH = 4.70
At 16.00 cm³, pH = 4.82
Interpolating the midpoint (15.00 cm³) gives approximately pH 4.76.
At half-neutralisation, pH = pKₐ.
pKₐ = 4.76
Kₐ = 10-pKₐ = 10-4.76 = 1.74 × 10⁻⁵ mol dm⁻³
pH of acid (at 0 volume) = 2.49
[H⁺] = 10-2.49 = 0.00324 mol dm⁻³
Kₐ = [H⁺]² / [HA]
[HA] = [H⁺]² / Kₐ
[HA] = (0.00324)² / 1.74×10⁻⁵
[HA] = 0.602 mol dm⁻³
Add the NaOH in smaller increments near the equivalence point.
This allows for a more precise determination of the vertical section and the equivalence volume.
- Set up a series of buffers with known, varying pHs.
- Place the pH meter in each solution, measuring the pH reading, and washing with distilled water between each solution.
- Plot a graph with known buffer pHs on the x-axis and the pH meter reading on the y-axis. Draw a line of best fit – this is the calibration curve.
- When using the pH meter in the experiment, record the meter reading and use the calibration curve to determine the true pH reading.