Acid-base titrations Revision Notes

An acid-base titration is a quantitative analytical technique used to determine the concentration of an acid or base solution. It involves the gradual addition of a solution of known concentration (the titrant) to a measured volume of another solution of unknown concentration (the analyte) until the reaction between them is complete. The point at which the reaction is complete is called the endpoint, which is detected by a suitable indicator that changes colour.

The titration relies on precise measurement of volumes and careful observation of colour changes to calculate the unknown concentration accurately. The main apparatus used are the burette, volumetric pipette, and indicator.

• Titrant: solution of known concentration, usually in the burette.
• Analyte: solution of unknown concentration, measured by volumetric pipette.
• Indicator: chemical that changes colour at the endpoint of the reaction.

The burette is a long, graduated glass tube with a tap at the bottom, designed to deliver variable volumes of liquid accurately, typically to the nearest 0.05 cm³. It is used to add the titrant solution to the analyte gradually and precisely.

Before starting, the burette is rinsed with the titrant solution to avoid dilution. It is then filled with the titrant, ensuring no air bubbles are present in the nozzle. The initial volume is recorded to two decimal places.

During titration, the tap is opened to allow the titrant to flow dropwise near the endpoint to avoid overshooting. The final volume is recorded, and the volume of titrant used is calculated by subtracting the initial reading from the final reading.

• Graduated scale allows volume measurement to ±0.05 cm³.
• Tap controls flow rate for precise delivery.
• Used to add titrant until the endpoint is reached.

The volumetric pipette is a piece of glassware designed to measure a fixed, accurate volume of liquid, commonly 25.0 cm³, with an uncertainty of ±0.03 cm³. It is used to transfer a known volume of the analyte solution into a conical flask for titration.

Before use, the pipette is rinsed with the analyte solution to prevent contamination. The liquid is drawn up to the calibration mark using a pipette filler, ensuring the bottom of the meniscus aligns exactly with the mark at eye level.

The analyte is then delivered into the conical flask by touching the tip to the inside of the flask and allowing the liquid to drain freely without blowing out the last drop, ensuring accurate volume delivery.

• Measures fixed volume accurately (e.g., 25.0 cm³).
• Rinsed with analyte to avoid dilution.
• Delivers analyte to conical flask for titration.

An indicator is a chemical that changes colour at (or very close to) the equivalence point of the titration, signalling that the acid and base have reacted completely. The choice of indicator depends on the strength of the acid and base involved, as different indicators change colour at different pH ranges.

For example, phenolphthalein changes from colourless to pink around pH 8.2–10, suitable for titrations involving strong base and weak acid. Methyl orange changes from red to yellow around pH 3.1–4.4, suitable for strong acid and weak base titrations.

The indicator is added in a few drops to the analyte solution before titration begins. The endpoint is identified by the first permanent colour change that persists when titrant is added dropwise.

• Indicator signals endpoint by colour change.
• Choice depends on acid-base strength and pH at equivalence point.
• Added to analyte before titration starts.

1. Rinse the burette with the titrant solution and fill it, recording the initial volume.

2. Rinse the volumetric pipette with the analyte solution, then use it to transfer a measured volume (e.g., 25.0 cm³) of analyte into a clean conical flask.

3. Add 2–3 drops of a suitable indicator to the analyte in the conical flask and swirl gently to mix.

• 4. Place the conical flask on a white tile to observe colour changes clearly.
• 5. Slowly open the burette tap to add titrant to the analyte, swirling the flask continuously.
• 6. Near the endpoint, add titrant dropwise until the indicator changes colour permanently.
• 7. Record the final burette reading and calculate the volume of titrant used.
• 8. Repeat the titration until at least two concordant titres (within 0.10 cm³) are obtained for accuracy.

The concentration of the unknown solution is calculated using the titration data and the balanced chemical equation. The general formula is:

\[ \text{Concentration of analyte (mol/dm}^3) = \frac{\text{moles of titrant}}{\text{volume of analyte (dm}^3)} \]

Moles of titrant are calculated from its concentration and volume used: \( \text{moles} = \text{concentration} \times \text{volume} \). Volumes must be converted from cm³ to dm³ by dividing by 1000.

• Balanced equation used to find mole ratio between titrant and analyte.
• Calculate moles of titrant from volume and concentration.
• Calculate moles of analyte using mole ratio.
• Calculate concentration of analyte from moles and volume.

An indicator is a chemical substance that changes colour at a particular pH range. In acid-base titrations, indicators are used to signal the end-point — the point at which the amount of titrant added is chemically equivalent to the amount of analyte present.

Indicators are weak acids or bases that have different colours in their protonated and deprotonated forms. The colour change occurs over a specific pH range, known as the indicator’s transition interval.

For example, methyl orange changes from red in acidic solution (pH < 3.1) to yellow in alkaline solution (pH > 4.4). Phenolphthalein is colourless in acidic and neutral solutions but turns pink in alkaline solutions (pH > 8.2).

• Indicator changes colour at a specific pH range.
• Colour change signals the end-point of the titration.
• Choice of indicator depends on the expected pH at equivalence point.

The end-point is identified by adding the titrant slowly to the analyte solution containing a few drops of indicator. The titrant is added from a burette while swirling the conical flask to mix thoroughly.

As the end-point approaches, the indicator’s colour will begin to change. The titrant must be added dropwise near the expected end-point to avoid overshooting.

The end-point is reached when a permanent colour change is observed — that is, the colour remains after swirling and does not revert on standing. This indicates that the reaction between acid and base is complete.

• Add titrant gradually, especially near the end-point.
• Swirl the solution continuously for even mixing.
• Stop adding titrant when the indicator shows a permanent colour change.

The choice of indicator depends on the type of acid-base reaction and the pH at the equivalence point.

For strong acid–strong base titrations, the equivalence point is around pH 7, so indicators like methyl orange or bromothymol blue (which changes colour near neutral pH) are suitable.

For strong acid–weak base titrations, the equivalence point is below pH 7, so methyl orange is preferred because it changes colour in acidic conditions (pH 3.1–4.4). For weak acid–strong base titrations, the equivalence point is above pH 7, so phenolphthalein (pH 8.2–10) is appropriate.

• Match indicator pH range to expected equivalence point.
• Use methyl orange for strong acid–weak base titrations.
• Use phenolphthalein for weak acid–strong base titrations.
• Use bromothymol blue for strong acid–strong base titrations.

Apparatus needed: burette, stand and clamp, conical flask, pipette, white tile, indicator, acid and base solutions of known concentration.

Method: 1. Rinse the burette with the titrant solution and fill it, noting the initial volume. 2. Use a pipette to transfer a measured volume of analyte to the conical flask. 3. Add 2–3 drops of a suitable indicator to the analyte. 4. Place the conical flask on a white tile to observe colour changes clearly. 5. Slowly add titrant from the burette while swirling the flask continuously. 6. Near the end-point, add titrant dropwise until a permanent colour change is observed. 7. Record the final burette reading and calculate the volume of titrant used.

Variables: independent variable is volume of titrant added; dependent variable is the colour change of the indicator; control variables include concentration and volume of analyte, temperature, and indicator type and amount. Repeat the titration at least three times to obtain concordant results (within 0.1 cm³). Calculate the average volume for accuracy and precision. Safety precautions include wearing eye protection and handling acids and bases carefully to avoid spills and skin contact.

• Use precise volumetric apparatus for accuracy.
• Add titrant slowly near end-point to avoid overshoot.
• Repeat titrations to obtain consistent results.
• Wear safety goggles and handle chemicals carefully.

1. Write the balanced equation: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

2. The equivalence point is at pH 7, but since phenolphthalein changes colour at pH 8.2–10, it is suitable for strong acid–strong base titrations.

3. Suppose the burette readings are initial 0.00 cm³ and final 24.80 cm³. Volume of NaOH used = 24.80 cm³ = 0.02480 dm³. 4. Calculate moles of NaOH added: concentration × volume = 0.100 mol/dm³ × 0.02480 dm³ = 0.00248 mol. 5. From the equation, moles HCl = moles NaOH = 0.00248 mol. 6. Calculate concentration of HCl: moles / volume = 0.00248 mol / 0.0250 dm³ = 0.0992 mol/dm³, close to the expected 0.100 mol/dm³.

• Balanced equation confirms 1:1 mole ratio.
• Indicator chosen matches expected pH range.
• Volume and concentration calculations confirm titration accuracy.