Enthalpy of dissolution of copper sulphate or potassium nitrate

Introduction

The enthalpy of dissolution represents a fundamental thermodynamic property that measures the heat change when one mole of a solute dissolves in a solvent to form an infinitely dilute solution. This experiment provides students with hands-on experience in calorimetry and thermodynamics while determining the heat of solution for common salts like copper sulphate (CuSO₄) and potassium nitrate (KNO₃).

Understanding dissolution enthalpy is crucial for various industrial applications, including pharmaceutical formulation, chemical processing, and environmental studies. This comprehensive guide will walk you through the complete experimental procedure for determining the enthalpy of dissolution in laboratory conditions.

Aim

To determine the enthalpy of dissolution of copper sulphate (CuSO₄) or potassium nitrate (KNO₃) using calorimetric methods in a controlled laboratory environment.

Keywords: enthalpy of dissolution, heat of solution, calorimetry experiment, copper sulphate enthalpy, potassium nitrate dissolution

Apparatus Required

Essential Equipment:

• Calorimeter (coffee cup or bomb calorimeter)
• Thermometer (0.1°C precision)
• Stirrer (glass rod or magnetic stirrer)
• Weighing balance (0.01g accuracy)
• Measuring cylinder (100ml)
• Beaker (250ml)
• Funnel
• Stopwatch
• Insulation material (cotton or cork)

Chemicals:

• Anhydrous copper sulphate (CuSO₄) or potassium nitrate (KNO₃)
• Distilled water

Theory

Definition of Enthalpy of Dissolution

The enthalpy of dissolution (ΔH_solution) is defined as the amount of heat absorbed or released when one mole of a solute dissolves completely in a solvent under constant pressure and temperature conditions.

Thermodynamic Principle

When a solute dissolves in a solvent, two primary processes occur:

1. Lattice Energy: Energy required to break ionic bonds in the solid crystal (endothermic)
2. Hydration Energy: Energy released when ions get surrounded by water molecules (exothermic)

The net enthalpy change is expressed as: ΔH_solution = ΔH_lattice + ΔH_hydration

Mathematical Expression

ΔH_solution = (m × s × ΔT) / n

Where:

• m = mass of solution (g)
• s = specific heat capacity of solution (4.18 J/g°C for dilute aqueous solutions)
• ΔT = temperature change (°C)
• n = number of moles of solute

Procedure

Preparation:

1. Clean and dry the calorimeter thoroughly
2. Record the mass of empty calorimeter
3. Add 100ml of distilled water to the calorimeter
4. Record the mass of calorimeter with water
5. Insert thermometer and stirrer through the lid

Experimental Steps:

Step 1: Initial Temperature Measurement

• Stir the water gently for 2-3 minutes
• Record the initial temperature (T₁) at 30-second intervals for 2 minutes

Step 2: Addition of Solute

• Weigh exactly 5g of CuSO₄ or KNO₃
• Quickly add the solid through the opening in the calorimeter lid
• Start timing immediately and stir continuously

Step 3: Temperature Monitoring

• Record temperature at 30-second intervals for 5 minutes
• Continue until temperature becomes constant or starts to decrease
• Note the maximum temperature reached (T₂)

Step 4: Data Collection

• Record final mass of calorimeter and solution
• Calculate moles of solute added

Observation Table

Raw Data Collection:

Calculated Data:

Result

Sample Calculation (Copper Sulphate):

Given:

• Mass of CuSO₄ = 5.00 g
• Moles of CuSO₄ = 5.00g ÷ 159.61 g/mol = 0.0313 mol
• Mass of solution = 105 g
• ΔT = 4.5°C
• Specific heat capacity (s) = 4.18 J/g°C

Calculation: q = m × s × ΔT = 105 × 4.18 × 4.5 = 1975.05 J ΔH = q/n = 1975.05 J ÷ 0.0313 mol = 63.1 kJ/mol

Expected Results:

Copper Sulphate (CuSO₄):

• Exothermic dissolution
• ΔH_solution ≈ -66.5 kJ/mol

Potassium Nitrate (KNO₃):

• Endothermic dissolution
• ΔH_solution ≈ +34.9 kJ/mol

Note: The negative sign indicates exothermic process, positive indicates endothermic.

Precautions

Chemical Handling:

1. Handle anhydrous CuSO₄ with care - it's toxic and can cause skin irritation
2. Avoid direct contact with chemicals; wear gloves and safety goggles
3. Wash hands immediately after handling chemicals
4. Avoid inhaling powdered chemicals

Experimental Precautions:

1. Ensure calorimeter is perfectly insulated to minimize heat loss
2. Add solute quickly to minimize heat exchange with surroundings
3. Stir solution gently but continuously to ensure uniform temperature
4. Record temperature immediately and at regular intervals
5. Use dry apparatus to prevent errors in mass measurement
6. Calibrate thermometer before use

Data Collection:

1. Take temperature readings at eye level to avoid parallax error
2. Wait for temperature stabilization before initial readings
3. Record all observations promptly and accurately

Viva Voce Questions and Answers

Q1: What is meant by enthalpy of dissolution?

A1: Enthalpy of dissolution is the heat change when one mole of a solute dissolves in a solvent to form an infinitely dilute solution under standard conditions.

Q2: Why is a polystyrene cup used as calorimeter?

A2: Polystyrene is a good thermal insulator with low heat capacity, minimizing heat loss to surroundings and providing accurate results.

Q3: What factors affect the enthalpy of dissolution?

A3: Factors include:

• Nature of solute and solvent
• Temperature
• Pressure
• Concentration of solution
• Ionic strength

Q4: Why should the solute be added quickly?

A4: Quick addition prevents heat loss to the surroundings and ensures accurate measurement of temperature change during dissolution.

Q5: What is the significance of stirring the solution?

A5: Stirring ensures uniform temperature distribution throughout the solution and promotes complete dissolution of the solute.

Q6: How would your result be affected if some solution is spilled during transfer?

A6: Results would be inaccurate as the actual amount of solute dissolved would be less than the measured amount, leading to incorrect enthalpy calculation.

Q7: What is the difference between endothermic and exothermic dissolution?

A7: Endothermic dissolution absorbs heat (ΔH positive) while exothermic dissolution releases heat (ΔH negative).

Q8: Why do we use distilled water in this experiment?

A8: Distilled water is free from impurities that could affect the dissolution process and introduce errors in heat measurements.

Applications and Significance

Industrial Applications:

• Pharmaceutical Industry: Understanding drug solubility and formulation
• Chemical Manufacturing: Process optimization and safety considerations
• Food Industry: Flavor compound dissolution studies
• Environmental Science: Pollutant dissolution in water bodies

Educational Value:

This experiment demonstrates fundamental principles of:

• Thermochemistry
• Solution chemistry
• Calorimetry techniques
• Data analysis and error calculation

Conclusion

The enthalpy of dissolution experiment provides valuable insights into the thermodynamic behavior of solutes in solution. Through careful measurement and analysis, students can understand the energy relationships in dissolution processes and develop essential laboratory skills in calorimetry.

Whether studying the exothermic dissolution of copper sulphate or the endothermic process of potassium nitrate, this experiment reinforces theoretical concepts while introducing practical applications in chemistry and related fields.