This practical guide outlines how to determine the enthalpy change of a reaction in aqueous solution (such as a neutralisation or displacement reaction) using a simple coffee-cup calorimeter. Accurate temperature readings over time and graphical extrapolation are used to correct for heat loss to the surroundings.
🔑 Core Specification Link
This practical is directly linked to topic 3.1.4 Energetics, focusing on the relation \( q = mc\Delta T \), standard enthalpy definitions, and Hess's law calculations.
Aim
To measure the temperature change over time for the reaction between hydrochloric acid and sodium hydroxide solution, and to calculate the standard enthalpy change of neutralisation.
Equipment List
- Expanded polystyrene cup (acts as the calorimeter)
- Plastic lid with a hole for the thermometer
- Glass beaker (250 cm³, to support the polystyrene cup)
- Thermometer (graduated to 0.1 °C or 0.2 °C) or temperature probe
- Measuring cylinders (50 cm³)
- Stirring rod (glass)
- Stopwatch or timer
Experimental Method
- Place the polystyrene cup inside the 250 cm³ glass beaker to provide stability and additional insulation.
- Using a measuring cylinder, transfer exactly 25.0 cm³ of 1.00 mol dm⁻¹ hydrochloric acid (\(\text{HCl}\)) into the polystyrene cup.
- Place the thermometer into the acid. Start the stopwatch and record the temperature of the acid every 30 seconds for 2 minutes to establish a baseline.
- Measure 25.0 cm³ of 1.00 mol dm⁻¹ sodium hydroxide solution (\(\text{NaOH}\)) in a second measuring cylinder.
- At exactly 2 minutes and 30 seconds, do not record the temperature. Instead, quickly add the sodium hydroxide solution to the polystyrene cup.
- Immediately place the lid on the cup, insert the thermometer, and stir the reaction mixture gently.
- Record the temperature of the solution every 30 seconds starting from 3 minutes, up to 7 minutes.
- Plot a graph of temperature (y-axis) against time (x-axis). Draw two lines of best fit: one through the temperature readings taken prior to mixing, and one through the readings taken during cooling after mixing.
- Extrapolate both lines to the time of mixing (2.5 minutes) to determine the theoretical maximum temperature change (\(\Delta T\)) if no heat loss had occurred.
Safety & Risk Assessment
| Hazard | Risk | Precaution |
|---|---|---|
| 1.00 mol dm⁻¹ Hydrochloric acid | Skin and eye irritation. | Wear safety goggles and a lab coat. Wash any skin splashes immediately. |
| 1.00 mol dm⁻¹ Sodium hydroxide | Skin irritation and potential eye damage. | Wear safety goggles. Wash skin spills immediately with water. |
| Glassware and thermometer | Cuts from broken glass. | Handle with care. Avoid using the thermometer as a stirring rod. Place in secure positions. |
Results & Calculations
Temperature-Time Extrapolation Graph
Because heat loss starts as soon as the reaction begins, the temperature change cannot be read simply as the difference between initial and maximum temperatures. Instead, we plot the temperature against time and extrapolate back to the time of mixing (2.5 minutes) as shown below:
Assume specific heat capacity \(c = 4.18\text{ J g}^{-1}\text{ K}^{-1}\) and density of solution \(\rho = 1.00\text{ g cm}^{-3}\).
Step 1: Calculate the total mass of the solution (\(m\))
\[ \text{Total volume} = 25.0 + 25.0 = 50.0\text{ cm}^3 \] \[ \text{Mass } m = \text{volume} \times \text{density} = 50.0\text{ cm}^3 \times 1.00\text{ g cm}^{-3} = 50.0\text{ g} \]Step 2: Calculate the heat energy transferred (\(q\))
\[ q = m \times c \times \Delta T = 50.0\text{ g} \times 4.18\text{ J g}^{-1}\text{ K}^{-1} \times 6.80\text{ K} = 1421.2\text{ J} = 1.4212\text{ kJ} \]Step 3: Calculate the moles of limiting reagent used (\(n\))
\[ n(\text{HCl}) = \text{concentration} \times \text{volume (dm}^3) = 1.00\text{ mol dm}^{-3} \times 0.0250\text{ dm}^3 = 0.0250\text{ mol} \] \[ n(\text{NaOH}) = 1.00\text{ mol dm}^{-3} \times 0.0250\text{ dm}^3 = 0.0250\text{ mol} \]Since the reaction ratio is 1:1, 0.0250 moles of water are formed.
Step 4: Calculate the standard molar enthalpy change (\(\Delta H\))
\[ \Delta H_{\text{neut}} = -\frac{q}{n} = -\frac{1.4212\text{ kJ}}{0.0250\text{ mol}} = -56.8\text{ kJ mol}^{-1} \]The standard enthalpy change of neutralisation is -56.8 kJ mol⁻¹ (3 significant figures).
Sources of Error & Improvements
| Error Source | Classification | Consequence & Mitigation |
|---|---|---|
| Heat loss to the surroundings | Systematic | Heat escapes through the cup walls and lid, reducing the maximum recorded temperature and making \(\Delta H\) less exothermic. Mitigation: Use a double-insulated polystyrene cup, add a lid, or use a vacuum flask. |
| Heat capacity of the cup and thermometer | Systematic | A small amount of heat is absorbed by the cup and the thermometer rather than the water. Mitigation: Account for the calorimeter constant (heat capacity of calorimeter) in calculations. |
| Uncertainty in thermometer reading | Random | Reading errors on analogue thermometers introduce uncertainty. Mitigation: Use a digital temperature probe connected to a datalogger. |
| Assuming density and specific heat capacity of solution are equal to water | Systematic | Solutes modify specific heat capacity and density. Mitigation: Use precise literature specific heat capacities for NaCl solutions. |
Common Exam Questions
1. Why is the temperature of the acid recorded for 2 minutes before adding the alkali?
This allows the acid solution to stabilise and reach thermal equilibrium with the surroundings, establishing an accurate baseline temperature for extrapolation.
2. Why does the extrapolation method provide a more accurate temperature change than simply subtracting the initial temperature from the maximum temperature?
Heat loss occurs continuously as the reaction proceeds. The maximum recorded temperature is always lower than the theoretical maximum temperature. Extrapolating the cooling curve back to the time of mixing corrects for the heat lost during the initial stages of the reaction.
3. A student reacted anhydrous copper(II) sulfate with water. State two key assumptions made when calculating the enthalpy change of the solution.
First, that the density of the final copper(II) sulfate solution is 1.00 g cm⁻³. Second, that the specific heat capacity of the solution is equal to that of pure water (4.18 J g⁻¹ K⁻¹).
CPAC Skills Assessed
- CPAC 2: Design and apply investigative approaches to measure thermal energy changes.
- CPAC 3: Safely use chemical substances and basic calorimetry apparatus.
- CPAC 4: Make and record temperature measurements over time with appropriate precision.
When calculating \(q = mc\Delta T\), ensure that \(m\) is the mass of the solution being heated (in this neutralisation reaction, the sum of both solution volumes: 50.0 g), not the moles or the mass of any dry reagent. Do not forget to include the negative sign in the final answer if the temperature increased (exothermic reaction).