Introduction: Measuring Enthalpy Changes

Welcome to Required Practical 2 (RP2)! In this chapter, we explore how to measure the "heat energy" exchanged during a chemical reaction. In Chemistry, we call this the enthalpy change (\(\Delta H\)). Whether you are watching a hand-warmer heat up or a cold pack chill down, you are witnessing an enthalpy change in action.

This practical is vital for your AQA A Level because it combines physical measurements (mass, volume, temperature) with mathematical calculations. Don't worry if the math seems daunting; we will break it down step-by-step.

Note: This practical supports your understanding of Section 3.1.4 (Energetics).

The Core Concept: Calorimetry

We measure enthalpy changes using a technique called calorimetry. Since we cannot measure "heat" directly, we measure the temperature change (\(\Delta T\)) of the surroundings (usually water or a solution) and use it to calculate the energy transferred.

The Fundamental Formula

To calculate the heat energy (\(q\)) transferred to the water, we use:

\(q = mc\Delta T\)

  • \(q\) = heat energy transferred (measured in Joules, \(J\)).
  • \(m\) = mass of the substance being heated/cooled (usually the water or solution in grams, \(g\)).
  • \(c\) = specific heat capacity (the energy needed to raise \(1 \text{ g}\) of a substance by \(1 \text{ K}\)). Note: You don't need to memorize this; it is usually provided as \(4.18 \text{ J g}^{-1} \text{ K}^{-1}\) for water.
  • \(\Delta T\) = the change in temperature (Kelvin, \(K\) or Celsius, \(^\circ\text{C}\)).

Quick Tip: In these experiments, we usually assume the density of an aqueous solution is the same as water (\(1 \text{ g cm}^{-3}\)). So, if you have \(50 \text{ cm}^3\) of solution, its mass is \(50 \text{ g}\).

The Experimental Setup: "The Coffee Cup" Calorimeter

For reactions involving solutions (like neutralisation or displacement), we use a simple polystyrene cup as our calorimeter. This is an example of an isobaric (constant pressure) calorimeter.

Why a Polystyrene Cup?

  • It is an excellent thermal insulator, which reduces heat loss to the surroundings.
  • It has a low heat capacity, meaning the cup itself absorbs very little of the reaction's heat.

The Procedure (Step-by-Step)

  1. Place a polystyrene cup inside a glass beaker for stability and extra insulation.
  2. Measure a known volume of a reagent (e.g., \(25 \text{ cm}^3\) of \(1.0 \text{ mol dm}^{-3} \text{ CuSO}_4\)) into the cup.
  3. Place a thermometer in the solution and record the temperature every minute for 3 minutes to establish a steady starting temperature.
  4. At the 4th minute, add the second reagent (e.g., zinc powder) but do not record the temperature yet. Stir the mixture thoroughly.
  5. Record the temperature every minute from the 5th minute until the 15th minute (or until the temperature has been falling steadily for several minutes).

Handling the Data: Extrapolation Graphs

One of the biggest challenges in calorimetry is heat loss. As soon as the reaction starts producing heat, some of that heat escapes into the air. This means the highest temperature we record is actually lower than the true maximum temperature.

To fix this, we use a temperature-time graph and extrapolate:

  • Plot your temperature readings against time.
  • Draw a line of best fit through the temperatures recorded before the reagents were mixed (the initial temperature).
  • Draw a line of best fit through the cooling curve (the temperatures recorded after the maximum was reached).
  • Extend both lines back to the 4th minute (the moment of mixing).
  • The vertical distance between these two lines at the 4th minute is your corrected \(\Delta T\).

Did you know? This extrapolation method allows us to estimate the temperature change that would have occurred if the reaction happened instantly and no heat was lost!

Calculating Molar Enthalpy Change (\(\Delta H\))

Once you have your corrected \(\Delta T\), follow these three steps to find the molar enthalpy change:

Step 1: Calculate \(q\)

Use \(q = mc\Delta T\). Remember, \(m\) is the mass of the liquid, not the solid added. Your answer will be in Joules (\(J\)).

Step 2: Calculate the Moles (\(n\))

Find the number of moles of the limiting reactant.
For solutions: \(n = \text{concentration} \times \text{volume (in } \text{dm}^3)\).
For solids: \(n = \frac{\text{mass}}{\text{Ar or Mr}}\).

Step 3: Calculate \(\Delta H\)

\(\Delta H = \frac{-q}{n}\) (if the reaction is exothermic) or \(\Delta H = \frac{q}{n}\) (if endothermic).

Crucial Check: Divide by \(1000\) to convert your answer into \(\text{kJ mol}^{-1}\). Always include a + sign for endothermic reactions or a - sign for exothermic reactions.

Key Takeaway: If the temperature goes up, the reaction is exothermic, and \(\Delta H\) must be negative.

Sources of Error and Improvements

In the exam, you may be asked why your experimental value is different from the "Data Book" value. Common reasons include:

  • Heat Loss: Even with a lid and polystyrene, some heat escapes. Improvement: Use extra insulation or a vacuum flask.
  • Non-standard Conditions: Standard enthalpy changes are defined at \(100 \text{ kPa}\). If your lab pressure is different, the result will vary.
  • Heat Capacity of Apparatus: We often ignore the heat absorbed by the cup and the thermometer.
  • Incomplete Reaction: If the reactants don't fully react, the temperature change will be smaller than expected.
  • Specific Heat Capacity Assumption: We assume the solution has the same \(c\) as pure water (\(4.18 \text{ J g}^{-1} \text{ K}^{-1}\)).

Quick Review: Success Checklist

Common Mistake to Avoid: When calculating \(m\) for \(q = mc\Delta T\), only include the mass of the substance that is changing temperature. If you add \(5 \text{ g}\) of magnesium to \(50 \text{ g}\) of acid, many exam boards prefer you to use \(50 \text{ g}\) as the mass, as the specific heat capacity used is for the water/solution, not the metal.

  • Did you convert volume to mass? (\(1 \text{ cm}^3 = 1 \text{ g}\))
  • Did you use the extrapolation method on the graph to find \(\Delta T\)?
  • Did you convert \(J\) to \(kJ\)?
  • Did you check the sign (\(+\) or \(-\)) of your final answer?
  • Is your answer to an appropriate number of significant figures (usually the same as the least accurate measurement provided)?

To see how these enthalpy changes can be used to find unknown values, check out the notes on Hess's Law in Section 3.1.4.