Practical Investigation: Analysing the Concentration of an Unknown Acid or Base by Titration

Key Terms

• Standard Solution: A solution with a known concentration.

• Titrant: The solution of known concentration used in the titration (often in the burette).

• Analyte: The solution with unknown concentration placed in the conical flask.

• Titre: The volume of titrant required to reach the endpoint.

• Aliquot: A specific volume of the analyte transferred into the conical flask.

• Equivalence Point: The point at which the amount of titrant added completely reacts with the analyte.

• Endpoint: The point at which a noticeable change (e.g., a color change) occurs, indicating the equivalence point has been reached.

• Titration Error: The difference between the equivalence point and the endpoint.

Acid/Base Titrations - Equipment and Setup

To carry out an acid/base titration, you will need:

• Burette: To deliver the titrant solution into the conical flask.

• Conical Flask: To hold the analyte solution.

• Pipette: To measure a specific volume of the standard solution.

• Volumetric Flask: To prepare the standard solution.

• Indicator: To show the endpoint of the titration (e.g., phenolphthalein).

Important: Make sure to draw diagrams of the experimental setup in your exams, label all equipment properly, and draw straight lines using a ruler.

Method for Acid/Base Titration

Step 1: Selection of a Primary Standard

• A primary standard is a substance that is pure, stable, and can be used to prepare a solution with a known concentration.

• Choose a strong acid or base for titrations with weak acids or bases.

• Examples: 

  • Acidic Primary Standards: Oxalic acid, benzoic acid.

  • Basic Primary Standards: Sodium carbonate (Na₂CO₃), sodium bicarbonate (NaHCO₃).

Step 2: Preparation of the Standard Solution

1. Dry the solid primary standard (if required) in a drying oven or desiccator.

2. Weigh the required amount of primary standard.

3. Dissolve the weighed amount in a beaker with demineralized water.

4. Transfer the solution into a volumetric flask and fill it up to the graduation mark with demineralized water.

5. Stopper and invert the flask 20 times to ensure the solution is homogeneous.

Step 3: Selection of an Appropriate Indicator

• The indicator should change color at the pH that corresponds to the equivalence point of the titration. 

  • For acidic solutions: Methyl orange (pH range 3.1 – 4.4).

  • For neutral solutions: Bromothymol blue (pH range 6.0 – 7.6).

  • For basic solutions: Phenolphthalein (pH range 8.3 – 10.0).

Step 4: Rinsing of Glassware

• Rinse glassware with demineralized water to avoid contamination.

• Pipette: Rinse with the solution it will deliver.

• Burette: Rinse with the titrant solution.

• Conical Flask and Volumetric Flask: Rinse with demineralized water.

Step 5: Preparing the Solutions for Titration

1. Prepare the Conical Flask: Use a pipette to transfer 25 mL of the standard solution (e.g., sodium hydroxide) into the conical flask.

2. Prepare the Burette: Rinse with demineralized water first, then with the titrant solution (e.g., acetic acid). Fill the burette with the titrant and ensure no air bubbles are trapped in the burette.

Step 6: Performing the Titration

1. Add the titrant to the analyte drop by drop, swirling the flask continuously.

2. Watch for the endpoint where the color change becomes permanent.

3. Note the volume of titrant used to reach the endpoint.

Step 7: Calculation

Use the titration data to calculate the unknown concentration of the analyte.

Worked Example

1. Balanced Equation: The titration of sodium carbonate with nitric acid is given by the equation: 2HNO₃ (aq) + Na₂CO₃ (aq) → 2NaNO₃ (aq) + H₂O (l) + CO₂ (g)

2. Primary Standard: Mass of Na₂CO₃ = 20g Molar mass of Na₂CO₃ = 106g/mol Moles of Na₂CO₃ = 20g / 106g/mol = 0.1887 mol Standard concentration of Na₂CO₃ solution = moles / volume = 0.1887 mol / 0.25 L = 0.7547 M

3. Average Titre: Titration volumes (mL): 24.30, 23.60, 23.55, 23.60 (average = 23.60 mL or 0.02360 L).

4. Amount of Na₂CO₃ Reacted: Moles of Na₂CO₃ reacted = concentration × volume = 0.7547 M × 0.02360 L = 0.01780 mol

5. Concentration of HNO₃: From the balanced equation, the ratio of HNO₃ to Na₂CO₃ is 2:1. Moles of HNO₃ reacted = 2 × 0.01780 mol = 0.03560 mol Concentration of HNO₃ = moles / volume = 0.03560 mol / 0.025 L = 1.424 M

Thus, the concentration of the nitric acid solution is 1.4 M (rounded to two significant figures).

Evaluating Christine’s Titration Method

Experimental Result

Christine calculated the concentration of acetic acid as lower than the actual concentration of 0.0750 mol/L.

Analysis of Procedure:

1. Weighing of Sodium Hydroxide Pellets:

   • Sodium hydroxide is hygroscopic, meaning it draws in water from the air. This increases its mass and reduces the accuracy of its weighed mass, meaning that it is a poor primary standard

2. Indicator Use:

   • The choice of phenolphthalein as the indicator is suitable, but the endpoint detection (faint pink color) can be difficult to spot accurately. A more visible indicator or a pH meter could help improve accuracy.

3. Rinsing of Glassware:

   • It was correctly noted that the burette should be rinsed with the acetic acid solution, but the conical flask should be properly rinsed with the same solution to avoid dilution.

Improvements:

• Accuracy: Ensure all glassware is clean and properly rinsed with the solution it will contain to avoid contamination.

• Endpoint Detection: Use a more precise method (e.g., pH meter) to detect the endpoint.

• Repetition: Increase the number of titrations to improve reliability and reduce errors.

By improving these steps, Christine’s titration procedure could yield more accurate and reliable results.