Determination of Enthalpy Changes (OCR A Level Chemistry A): Revision Note

Exam code: H432

Richard Boole

Written by: Richard Boole

Reviewed by: Philippa Platt

Updated on

PAG 3: Determination of enthalpy changes

  • Calorimetry is a technique used to measure changes in enthalpy of chemical reactions

  • calorimeter can be a polystyrene drinking cup, a vacuum flask or metal can

Thermal measurement setup with a thermometer in a polystyrene cup containing a reaction mixture, topped with a plastic lid.
A polystyrene cup can act as a calorimeter to find enthalpy changes in a chemical reaction
  • Specific heat capacity is the energy needed to raise the temperature of 1 g of a substance by 1 K

  • The specific heat capacity of water is 4.18 J g-1 K-1

  • The energy transferred as heat can be calculated by:

q = m c ΔT

  • Where:

    • q = heat energy transferred, J

    • m = mass of water, g

    • c = specific heat capacity, J g-1 K-1

    • ΔT = temperature change, K

  • There are two types of calorimetry experiments for you to know:

    • Enthalpy changes of reactions in solution

    • Enthalpy changes of combustion

  • For both calorimetry experiments, you should be able to:

    • Outline the experiment

    • Process experimental data using calculations or graphical methods

PAG 3.1: Determination of the enthalpy change of neutralisation

  • The principle of these calorimetry experiments is to mix stoichiometric quantities of acid and alkali then measure the temperature change over the course of a few minutes

  • The apparatus needed to carry out an enthalpy of reaction in solution calorimetry experiment is shown above

Sample method for a neutralisation reaction

  1. Using a measuring cylinder, place 25.0 cm3 of 1.0 mol dm-3 hydrochloric acid into the polystyrene cup

  2. Using a measuring cylinder, place 25.0 cm of 1.0 mol dm-3 sodium hydroxide solution into a separate polystyrene cup 

  3. Draw a table to record:

    • The initial temperature of the HCl (aq) and NaOH (aq)

    • The temperature of the combined reaction mixture every half minute up to 9.5 minutes

  4. Put a thermometer or temperature probe in each cup, stir, and record the temperature every half minute for 2½ minutes

  5. At precisely 3 minutes, add the contents of the two cups together

    • DO NOT RECORD THE TEMPERATURE AT 3 MINUTES

  6. Continue stirring and recording the temperature for 6 minutes

  • For the purposes of the calculations, the following experimental assumptions are made:

    • That the specific heat capacity of the solution is the same as pure water, i.e. 4.18 J g-1 K-1

    • That the density of the solution is the same as pure water, i.e. 1 g cm-3

    • The specific heat capacity of the container is ignored

    • The reaction is complete

    • There are negligible heat losses

Temperature correction graphs

  • For reactions which are not instantaneous there may be a delay before the maximum temperature is reached

  • During that delay the substances may lose heat to the surroundings

    • This means that the true maximum temperature is never reached

  • To overcome this problem we can use graphical analysis to determine the maximum enthalpy change

Diagram of a heating experiment with labels: insulating lid, thermometer, draught shield, copper can with water, and spirit burner.

The steps to make a temperature correction graph are:

  1. Take a temperature reading before adding the reactants for a few minutes to get a steady value

  2. Add the second reactant and continue recording the temperature and time

  3. Plot the graph 

  4. Extrapolate the cooling section of the graph until you intersect the time at which the second reactant was added

  • For the neutralisation experiment, there can be two main possibilities:

    1. The initial temperatures of the HCl (aq) and NaOH (aq) are the same

      • They can both be plotted as the same line

      • The graph can be extrapolated to determine the change in temperature at 3 minutes

      • The subsequent q = m c ΔT calculation can be completed

    2. The initial temperatures of the HCl (aq) and NaOH (aq) are the different

      • The initial temperature of the HCl (aq) is plotted as one line

      • The initial temperature of the NaOH (aq) is plotted as a separate line on the same graph

      • The temperature of the reaction mixture can be plotted

      • Two extrapolations are performed

        • One for the HCl (aq) and another for the NaOH (aq)

      • The subsequent q = m c ΔT calculations for the HCl (aq) and NaOH (aq) are completed

      • Calculating the average of those two values gives the overall enthalpy change of neutralisation

PAG 3.2 Determination of an enthalpy change of reaction by Hess' law

  • This practical is an extended variation of PAG 3.1

Examiner Tips and Tricks

The exam board recommends:

  • The reaction of potassium carbonate with hydrochloric acid

  • The reaction of potassium hydrogen carbonate with hydrochloric acid

This allows the enthalpy change of reaction for the decomposition of potassium hydrogen carbonate to potassium carbonate to be determined

Diagram of enthalpy change: 2KHCO₃ to K₂CO₃, H₂O, CO₂; with HCl forms 2KCl, 2H₂O, 2CO₂. ΔHr = ΔH₁ - ΔH₂. Blue curve indicates reaction path.
  • Another common experiment for this Hess' law approach is hydrated copper sulfate and anhydrous copper sulfate to form copper sulfate solution

Energy cycle diagram showing transition from CuSO4 and water to hydrated CuSO4.5H2O. Arrow paths show enthalpy changes ΔHr, ΔH1, and ΔH2.

Sample method for a neutralisation reaction

  1. Using a measuring cylinder, place 25.0 cm3 of 1.0 mol dm-3 hydrochloric acid into the polystyrene cup

  2. Draw a table to record:

    • The initial temperature of the HCl (aq)

    • The temperature of the combined reaction mixture every half minute up to 9.5 minutes

  3. Using a top-pan balance, measure 3.25 g of potassium hydrogen carbonate , KHCO3 

  4. Put a thermometer or temperature probe in the cup of HCl (aq), stir, and record the temperature every half minute for 2½ minutes

  5. At precisely 3 minutes, add the potassium carbonate to the cup of HCl (aq)

    • DO NOT RECORD THE TEMPERATURE AT 3 MINUTES

  6. Continue stirring and record the temperature for 6 minutes

  7. Plot a temperature correction graph to determine the enthalpy change for the reaction

  8. Repeat this experiment using 2.50 g of potassium carbonate, K2CO3 

    • Both experiments can be repeated to find an average temperature changes

  9. Use the equation q = m c ΔT to calculate:

    • ΔH1 (enthalpy change of reaction for KHCO3 with HCl)

    • ΔH2 (enthalpy change of reaction for K2CO3 with HCl)

  10. Apply Hess' law to calculate ΔHr (enthalpy change for the decomposition of KHCO3 to K2CO3)

    • Using ΔHr = ΔH1 - ΔH2

PAG 3.3 Determination of Enthalpy Changes of Combustion

  • The principle here is to use the heat released by a combustion reaction to increase the heat content of water

  • A typical simple calorimeter is used to measure the temperature changes to the water

Simple Calorimeter_1, downloadable IB Chemistry revision notes

A simple combustion calorimeter

  • To complete this experiment, the following steps will need to be completed:

Three-step experiment: Step 1, measure water temperature in copper can; Step 2, weigh spirit burner with and without fuel; Step 3, heat and calculate fuel burnt.
  • It is important that you record:

    • The starting temperature

    • The final temperature

    • The starting mass of the spirit burner

    • The final mass of the spirit burner

  • Using this information, the mass of the fuel burned during the reaction and the temperature change can be deduced

  • This mass of fuel burned can be used to calculate the moles of fuel burned

  • The moles of fuel burned can be used to convert q to an enthalpy change

Key points to consider

  • Not all the heat produced by the combustion reaction is transferred to the water

    • Some heat is lost to the surroundings

    • Some heat is absorbed by the calorimeter

  • To minimise the heat losses the copper calorimeter should not be placed too far above the flame and a lid placed over the calorimeter

  • Shielding can be used to reduce draughts

  • In this experiment the main sources of error are

    • Heat losses

    • Incomplete combustion

Worked Example

1.023 g of propan-1-ol (M = 60.11 g mol-1) was burned in a spirit burner and used to heat 200 g of water in a copper calorimeter.

The temperature of the water rose by 30 oC.

Calculate the enthalpy of combustion of propan-1-ol using this data.

Answer:

  1. Calculate q

q = m c ΔT

q = 200 g x 4.18 J g–1 K–1 x 30 K = – 25 080 J

  1. Calculate the amount of propan-1-ol burned

moles = mass over M subscript straight r

moles = fraction numerator 1.023 space straight g over denominator 60.11 space straight g space mol to the power of negative 1 end exponent end fraction = 0.01702 mol

  1. Calculate ΔH

ΔH = q over n

ΔH = fraction numerator negative 25080 space straight J over denominator 0.01702 space mol end fraction = – 1 473 560 J

ΔH = -1 474 kJ = -1.5 x 103 kJ mol-1

Practical skills reminder

  • These calorimetry-based experiments develop a range of essential skills, including:

    • Accurately measuring temperature changes using a thermometer or probe

    • Measuring mass using a top-pan balance (e.g. for solids or spirit burners)

    • Handling volumes of liquids using measuring cylinders and polystyrene cups

    • Recording and plotting time–temperature data to apply graphical corrections

    • Calculating enthalpy changes using experimental data and the formula
      q = mcΔT

    • Applying Hess' law to derive enthalpy changes indirectly from two reactions

    • Identifying and explaining sources of experimental error, such as heat loss or incomplete combustion

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Richard Boole

Author: Richard Boole

Expertise: Chemistry Content Creator

Richard has taught Chemistry for over 15 years as well as working as a science tutor, examiner, content creator and author. He wasn’t the greatest at exams and only discovered how to revise in his final year at university. That knowledge made him want to help students learn how to revise, challenge them to think about what they actually know and hopefully succeed; so here he is, happily, at SME.

Philippa Platt

Reviewer: Philippa Platt

Expertise: Chemistry Content Creator

Philippa has worked as a GCSE and A level chemistry teacher and tutor for over thirteen years. She studied chemistry and sport science at Loughborough University graduating in 2007 having also completed her PGCE in science. Throughout her time as a teacher she was incharge of a boarding house for five years and coached many teams in a variety of sports. When not producing resources with the chemistry team, Philippa enjoys being active outside with her young family and is a very keen gardener