1. ENERGETICS CORE LEVEL

4. 5.2 Calculation of Enthalpy Changes (3h) 5.2.1 Calculate the heat change when the temperature of a pure substance is changed Students should be able to calculate the heat change for a substance given the mass, specific heat and temperature change. 5.2.2 Design suitable experimental procedures for measuring the heat energy change of a reaction. Students should consider in aqueous solution and combustion reactions. Use of the bomb calorimeter and calibration of calorimeters will not be assessed. 5.2.3 Calculate the enthalpy change for a reaction in aqueous solution using experimental data on temperature changes, quantities of reactants and mass of solution Enthalpy change of an acid–base reaction could be investigated. 5.2.4 Evaluate the results of experiments to determine enthalpy changes. Students should be aware of the assumptions made and errors due to heat loss.

5. Measuring Enthalpy Changes Enthalpy changes are measured in KJ per mole, but experiments on measuring enthalpy changes always measure a temperature change. The temperature change is then converted to enthalpy change using the following formula: Energy = m x c x change in temperature = mc  T m = mass in grams c = specific heat capacity of water in Joules/Kelvin/gram = 4.18 J /K / g N.B. 1:If you are carrying out an experiment where you heat up something other than water, you must use the specific heat capacity of that substance. 2. A 1K temperature change is the same as a 1 o C change If an aluminium block weighing 10 g is heated by a chemical reaction from 25 o C to 40 o C what is the amount of heat energy produced by this reaction? (The specific heat capacity of Al is 0.9 J per g per K)

6. Working Out Enthalpy Change from the Temperature Change 14 grams of sodium hydroxide pellets were dissolved in 100 cm3 of water. The temperature before adding the sodium hydroxide pellets was 25 degrees C, and after adding the pellets it was 35 degrees C. Calculate the enthalpy change in KJ/mole of the reaction (Specific heat capacity of water = 4.18 J/K/g) Step 1: Work out the energy change using the formula energy = m x c x change in temperature m = 100 grams (100 cm 3 of water is the “same” as 100 grams of water) change in temperature = 10 K energy = 100 x 4.18 x 10 J Step 2: Convert the enthalpy change to kJ by dividing by 1000 energy = 100 x 4.18 x 10 KJ 1000

7. Step 3: Work out how many moles there are …0.1 moles (moles = mass/formula mass of NaOH)) Step 4: Work out the enthalpy change per mole for this exothermic reaction so the enthalpy change for 1 mole = 4.2 / 0.1 = - 42 kJ/mol

8. Experiment to measure some energy changes of reactions occurring in solutions and to use the results to calculate enthalpy changes The experiment is done by simple calorimeter experiments which involve carrying out the reaction in a expanded polystyrene cup (a good insulator) and measuring the change in temperature. This is also a common practical assessment, and you should be aware of the sources of error in your measurements such as heat loss to surroundings, incorrect readings, inaccurate thermometers etc. Apparatus: Thermometer ? range and sensitivity? Lid + Expanded polystyrene cup used as a calorimeter Method: The reaction between copper (II) sulphate and zinc Caution: Zn dust is flammable 1. Write the balanced ionic equation for the reaction: 2. Measure out 25cc of 0.2M copper (II) sulphate solution into the calorimeter. 3. Measure the temperature of the solution. 4. Add 0.1mol of Zn. 5. Stir gently and continuously and note the highest temperature reached. 6. Work out the temperature change to the nearest 0.1 o C 7. Now try to work out: a) the energy given out by the reaction b) the enthalpy change per mol of Cu 2+ 8. Can you show the zinc was added in excess? 9. What key assumptions have you made in these calculations

13. 5.3 Hess’s Law (2h) 5.3.1 Determine the enthalpy change of a reaction which is the sum of two or more reactions with known enthalpy changes. Students should be able to use simple enthalpy cycles and enthalpy level diagrams and to manipulate equations. Students will not be required to state Hess’s law. 5.4 Bond Enthalpies (2h) 5.4.1 Define the term average bond enthalpy Bond enthalpies are quoted for the gaseous state and should be recognized as average values obtained from a number of similar compounds. 5.4.2 Explain, in terms of average bond enthalpies, why some reactions are exothermic and others are endothermic

14. 5.3 Hess’s Law (2h) What does Hess’ Law say? Why is Hess’s Law useful? What link is there between the enthalpy change of the forward and the back reaction? Now try to use Hess’ Law to answer the following questions: Using the chemical equations below: C(s) + O 2 (g)  CO 2 (g)  H = -390 kJ mol -1 Mn(s) + O 2 (g)  MnO 2 (s)  H = -520 kJ mol -1 Work out the enthalpy change for the following reaction: MnO 2 (s) + C(s)  Mn(s) + CO 2 (g)

15. Method 2 Using Equations

16. Using the chemical equations below: Cu(s) + ½ O 2 (g) –CuO(s)  H = -156 kJ mol -1 2Cu(s) + ½ O 2 (g)  Cu 2 O (s)  H = -170 kJ mol -1 work out the value of  H (in kJ) for the following reaction: 2CuO (s)  Cu 2 O (s) +½ O 2 (g)

17. Another way of using Hess’s Law

19. If we can measure ΔH2 and ΔH1, we can find ΔH. Using Hess's Law: ΔH + ΔH2 = ΔH1 Hence: ΔH = ΔH1 - ΔH2 Enthalpy cycles are useful because they enable a value for an enthalpy change to be determined for a reaction which cannot be determined directly from experiment.

25. Bonds and Enthalpy Cycles The bond-breaking and bond-making can be represented in an enthalpy cycle. Now can you use your data book and the information about bond enthalpies to calculate  H2 ?

27. Born-Haber cycle calculations

28. Via an Energy level cycle. This is the method always required in examinations The sum of energies for route 1 = sum of energies for route 2