Titration Basics (Higher Tier Overview)

Purpose of Titration

Titration is a laboratory method used to find the concentration of an unknown solution by reacting it with a solution of known concentration (called the standard solution or titrant). It relies on an acid-base neutralisation reaction where the acid and alkali react to form water and a salt.

An indicator is added to the unknown solution to show the end point of the titration. The end point is the moment when the acid and alkali have completely reacted, indicated by a permanent colour change.

The main goal is to measure accurately how much titrant is needed to neutralise the unknown solution, which then allows calculation of the unknown concentration.

Common indicators include phenolphthalein, which changes from colourless in acid to pink in alkali, and methyl orange, which changes from red in acid to yellow in alkali. These indicators are chosen because their colour change occurs close to the equivalence point of the titration.

Equipment and Setup

• Burette: A long, graduated glass tube with a tap at the bottom, used to add the titrant dropwise to the unknown solution. It allows precise measurement of the volume added.
• Pipette: Used to measure a fixed volume of the unknown solution accurately and transfer it to the conical flask.
• Conical Flask: The flask where the unknown solution and indicator are mixed. Its shape helps prevent splashing during swirling.
• White Tile: Placed under the conical flask to make the colour change of the indicator easier to see clearly.

Procedure Steps

1. Rinse the burette with the standard solution (titrant) and then fill it, ensuring no air bubbles are present. Record the initial volume.
2. Use a pipette to measure a precise volume of the unknown solution and transfer it to the conical flask.
3. Add a few drops of a suitable indicator to the unknown solution in the flask. Common indicators include phenolphthalein (pink in alkali, colourless in acid) or methyl orange (red in acid, yellow in alkali).
4. Slowly add the titrant from the burette to the flask, swirling the flask continuously to mix.
5. Stop adding titrant as soon as the indicator changes colour permanently, signalling the end point.
6. Record the final volume of titrant in the burette and calculate the volume used by subtracting the initial volume.

For example, if the initial burette reading is 0.00 cm³ and the final reading is 23.45 cm³, the volume of titrant used is:

$$\text{Volume of titrant} = 23.45 - 0.00 = 23.45 \text{ cm}^3$$

Calculations

The key to titration calculations is using the volumes measured and the balanced chemical equation to find the unknown concentration.

Steps for calculation:

• Read burette volumes accurately: Always record initial and final readings to find the volume of titrant used.
• Calculate the volume of titrant used: Subtract initial from final volume.
• Use the balanced chemical equation: Determine the mole ratio between the titrant and the unknown solution.
• Calculate moles of titrant used: Use the concentration and volume of the titrant.
• Calculate moles of unknown solution: Use the mole ratio from the balanced equation.
• Calculate the concentration of the unknown: Use the moles and volume of the unknown solution.

For example, if 25.0 cm³ of an acid is neutralised by 23.45 cm³ of 0.100 mol/dm³ sodium hydroxide (NaOH), the moles of NaOH used are:

$$\text{Moles of NaOH} = \text{concentration} \times \text{volume (dm}^3) = 0.100 \times \frac{23.45}{1000} = 0.002345 \text{ mol}$$

If the balanced equation is:

$$\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$$

The mole ratio of HCl to NaOH is 1:1, so moles of HCl = 0.002345 mol.

Then, the concentration of HCl is:

$$\text{Concentration of HCl} = \frac{\text{moles}}{\text{volume (dm}^3)} = \frac{0.002345}{25.0/1000} = 0.0938 \text{ mol/dm}^3$$

Example calculation inline: Calculate the moles of 0.100 mol/dm³ NaOH used if 25.0 cm³ is added during titration.

$$\text{Moles of NaOH} = 0.100 \times \frac{25.0}{1000} = 0.0025 \text{ mol}$$