During an acid-base neutralization, the pH change is non-linear. By tracking pH continuously as base is added, we obtain a pH titration curve. This guide explores the differences between strong and weak acid titration curves, indicator selection, and calculating the acid dissociation constant (\(K_{\text{a}}\)).
Core Specification Link
This practical supports 3.1.12 Acids and bases, focusing on pH curves, titration calculations, indicators, buffer action, and \(K_{\text{a}}\) determination.
Apparatus Setup
To record titration curves, a pH electrode is submerged in an acid solution on a magnetic stirrer while a base is added incrementally from a burette:
Aim
To investigate how pH changes when (a) a strong acid reacts with a strong base, (b) a weak acid reacts with a strong base, and (c) a strong acid reacts with a weak base (specifically HCl + NH₃), and to select suitable indicators based on the resulting pH titration curves.
Equipment List
- pH meter with electrode (calibrated)
- pH buffer solutions (pH 4.00, 7.00, and 10.00)
- 50 cm³ burette and stand
- 0.10 mol dm⁻³ hydrochloric acid (HCl)
- 0.10 mol dm⁻³ ethanoic acid (CH₃COOH)
- 0.10 mol dm⁻³ sodium hydroxide (NaOH)
- 0.10 mol dm⁻³ aqueous ammonia (NH₃)
- 250 cm³ beaker
- Magnetic stirrer and stirring bar
- Distilled water wash bottle
Experimental Method
Calibration of the pH Electrode
A pH meter measures the potential difference across a glass membrane. This potential shifts over time (electrode drift). To ensure accurate measurements, the meter must be calibrated using standard buffer solutions:
- Rinse the pH electrode thoroughly with distilled water and gently blot it dry.
- Submerge the electrode in the pH 7.00 buffer . Adjust the meter reading to read exactly 7.00.
- Rinse the electrode again and place it in the pH 4.00 buffer (or pH 10.00 buffer) . Adjust the scale calibration to match.
pH Measurements
- Pipette 25.0 cm³ of 0.10 mol dm⁻³ hydrochloric acid into a 250 cm³ beaker . Place the beaker on the magnetic stirrer, add the stirrer bar, and turn on the stirrer at a slow, steady speed.
- Position the pH electrode in the beaker so the bulb is fully submerged but clear of the rotating magnetic stirrer bar . Record the initial pH (at 0.0 cm³ base added).
- Fill the burette with 0.10 mol dm⁻³ sodium hydroxide solution.
- Add the sodium hydroxide in 1.0 cm³ portions . Record the volume added and the stable pH reading after each addition.
- Near the equivalence point (when pH begins changing rapidly, around pH 3 to 11): Reduce the addition increments to 0.20 cm³ to define the steep section of the curve.
- After passing the equivalence point and once the pH changes slow down again (above pH 11), return to 1.0 cm³ increments until a total of 40 cm³ of base has been added.
- Repeat steps 1 to 6 using 25.0 cm³ of 0.10 mol dm⁻³ ethanoic acid instead of hydrochloric acid (weak acid / strong base).
- Repeat the titration for a strong acid / weak base combination by titrating 25.0 cm³ of 0.10 mol dm⁻³ hydrochloric acid with 0.10 mol dm⁻³ aqueous ammonia in the burette.
Analysing pH Titration Curves
1. Strong Acid + Strong Base (HCl + NaOH)
- Starting pH: Very low (pH ≈ 1.00 for 0.10 mol dm⁻³ HCl).
- Shape: Remains flat and changes slowly until close to the equivalence point.
- Equivalence Point: Exactly pH 7.00. The vertical section is long and steep, stretching from pH 3 to 11.
- Indicator Choice: Either phenolphthalein or methyl orange is suitable because their transition ranges fall completely within the steep vertical section of the curve.
2. Weak Acid + Strong Base (CH₃COOH + NaOH)
- Starting pH: Higher than the strong acid (pH ≈ 3.00 because ethanoic acid is only partially ionised).
- Buffer Region: The pH rises slightly initially and then levels off into a buffer region (between 5 and 20 cm³ of base). The solution contains a mixture of unreacted weak acid (CH₃COOH) and its conjugate base (CH₃COO⁻), which resists changes in pH.
- Equivalence Point: Greater than 7.00 (typically pH 8.5 to 9.00). This occurs because of salt hydrolysis: \[ \text{CH}_3\text{COO}^-(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightleftharpoons \text{CH}_3\text{COOH}(\text{aq}) + \text{OH}^-(\text{aq}) \]
- Indicator Choice: Phenolphthalein is suitable (range 8.2 to 10.0). Methyl orange is unsuitable because its pH range (3.1 to 4.4) lies below the steep vertical section of the curve.
The ethanoate ions react with water to form hydroxide ions, making the neutralised solution basic.
3. Strong Acid + Weak Base (HCl + NH₃)
- Starting pH: Very low (pH ≈ 1.00 for 0.10 mol dm⁻³ HCl).
- Shape: Remains low and flat until close to the equivalence point. After the equivalence point, it levels off at pH ≈ 10 to 11 because aqueous ammonia is a weak base.
- Equivalence Point: Less than 7.00 (typically pH 5.0 to 5.5). This occurs because of salt hydrolysis of the ammonium ion: \[ \text{NH}_4^+(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightleftharpoons \text{NH}_3(\text{aq}) + \text{H}_3\text{O}^+(\text{aq}) \]
- Indicator Choice: Methyl orange is suitable (range 3.1 to 4.4) because the vertical section falls across the acidic region. Phenolphthalein is unsuitable because its range (8.2 to 10.0) lies well above the vertical section of the curve.
The ammonium ion acts as a weak acid in water, producing hydronium ions that make the neutralised ammonium chloride solution weakly acidic.
Step 1: Identify the half-equivalence point
\[ \text{Half-equivalence volume} = \frac{\text{Equivalence volume}}{2} = \frac{24.80}{2} = 12.40\text{ cm}^3 \]Step 2: Relate pH and \(\text{p}K_{\text{a}}\)
At the half-equivalence point, exactly half of the weak acid (HA) has been converted to its conjugate base (A⁻), meaning:
\[ [\text{HA}] = [\text{A}^-] \]Substituting this equality into the Henderson-Hasselbalch equation:
\[ \text{pH} = \text{p}K_{\text{a}} + \log\left(\frac{[\text{A}^-]}{[\text{HA}]}\right) \implies \text{pH} = \text{p}K_{\text{a}} + \log(1) \implies \text{pH} = \text{p}K_{\text{a}} \]Therefore, at the half-equivalence point:
\[ \text{p}K_{\text{a}} = \text{pH} = 4.76 \]Step 3: Convert \(\text{p}K_{\text{a}}\) to \(K_{\text{a}}\)
\[ K_{\text{a}} = 10^{-\text{p}K_{\text{a}}} = 10^{-4.76} = 1.74 \times 10^{-5}\text{ mol dm}^{-3} \]The acid dissociation constant of the weak acid is 1.74 × 10⁻⁵ mol dm⁻³.
Safety & Risk Assessment
| Hazard | Risk | Precaution |
|---|---|---|
| 0.10 mol dm⁻³ NaOH | Skin irritant; corrosive to eyes. | Wear safety goggles and lab coat. Wash splashes off skin immediately. |
| 0.10 mol dm⁻³ ethanoic acid | Pungent vapor, skin and eye irritant. | Use in a well-ventilated laboratory; wear safety goggles. |
| Glassware & pH Electrode | Breakage can cause cuts. Glass bulb on probe is fragile. | Handle electrode carefully. Avoid hitting the probe with the rotating magnetic stirrer bar. |
Sources of Error & Improvements
- Temperature variations: pH measurements are temperature-dependent. If the temperature changes during the titration, the curves will skew. Improvement: Keep solutions in a constant-temperature water bath and calibrate the pH meter at the experimental temperature.
- Slow response time of the electrode: The pH meter takes a short time to equilibrate after base addition. Recording the pH too quickly leads to lag in the steep section. Improvement: Wait 10 to 15 seconds after each addition before recording the pH.
Exam questions, marked instantly
Write your answer and the AI examiner checks it against the mark scheme, showing exactly which marks you earned and what to add.
Exam-style question
2 marksExplain why the vertical section of a weak acid–strong base titration curve is shorter than that of a strong acid–strong base curve.
Show Mark Scheme
- A weak acid is only partially dissociated, so the pH is higher before the end point (a buffer region forms) [1]
- So the steep section starts at a higher pH, about 7 rather than about 3, and is shorter [1]
Exam-style question
2 marksExplain why no indicator is suitable for a titration between a weak acid and a weak base.
Show Mark Scheme
- There is no steep (vertical) section at the equivalence point; the pH changes gradually [1]
- An indicator changes colour over about 2 pH units, so the colour would change gradually over a large volume and there would be no sharp end point [1]
CPAC Skills Assessed
- CPAC 2: Calibrates and uses a digital pH meter to monitor a reaction.
- CPAC 3: Safely sets up burettes and magnetic stirrers.
- CPAC 4: Records and plots pH values at varying volume increments.
When selecting a suitable indicator, always state that the pH range of the indicator’s colour change must fall completely within the vertical section of the pH curve.