Edexcel International A Level Chemistry

Topic 6: Energetics

Chemical reactions move heat between the chemicals and their surroundings. This topic gives that heat a name, a sign and a number, shows how to measure it, and shows how to work it out when it cannot be measured.

The one idea

A reaction breaks bonds in the reactants, which takes in energy, and makes bonds in the products, which gives energy out. The balance is the enthalpy change, ΔH. If heat leaves the chemicals, ΔH is negative and the reaction is exothermic. If heat enters them, ΔH is positive and it is endothermic.

Almost every question in this topic is one of three moves:

1. Name and sign it. Know which of the five enthalpy changes you are being asked for, write its equation with state symbols, and give ΔH the right sign and units.
2. Measure it. Turn a thermometer reading into kJ mol⁻¹ with q = mcΔT, then judge the answer: errors, uncertainty, assumptions and a cooling-curve correction.
3. Calculate it without measuring it. Use Hess's Law with formation or combustion data, or add up bond enthalpies.

Where to find things

How to use these notes

• Read the pages in order the first time. Pages 1 to 4 give you the language. Pages 5 to 8 use it to measure. Pages 9 to 12 use it to calculate indirectly. Pages 13 to 15 use bonds.
• Each page opens with its one idea in a sentence. If you can explain it to someone else, you know the page.
• Coloured boxes mean the same thing on every page: definitions you may be asked to quote, key ideas, common mistakes, exam technique, and beyond the spec (interesting, but you will not be examined on it).
• Each page ends with questions and fully worked answers. Cover the answers, try the questions, then mark yourself.
• Use the topic on a page the night before an exam.

This topic has Core Practical 2, the determination of the enthalpy change of a reaction using Hess's Law, on pages 11 and 12. The data in the practical pages are illustrative: they show a typical set of results, not measured values.

Constants used in this topic: specific heat capacity of water and dilute solutions 4.18 J g⁻¹ °C⁻¹, density of solutions 1.00 g cm⁻³, standard conditions 100 kPa and 298 K. Relative atomic masses are the Edexcel IAL values to one decimal place: H 1.0, C 12.0, N 14.0, O 16.0, Mg 24.3, Cl 35.5.

The skills that earn most marks here are writing the exact equation for each named enthalpy change; keeping the sign and the units on every ΔH; using the correct mass (water or solution, never the reactant) in q = mcΔT; drawing the cycle and following the arrows in a Hess calculation; and counting bonds from the equation, not from one molecule.

Pages

1. Enthalpy change and its sign

   ΔH is the heat energy change at constant pressure, measured under standard conditions. A negative ΔH means heat leaves the chemicals (exothermic); a positive ΔH means heat enters them (endothermic).

   11 min 6.16.2

2. Enthalpy level diagrams

   An enthalpy level diagram is a picture of ΔH. Reactants and products sit on two levels, and an arrow from reactants to products carries the sign and the size of the change.

   10 min 6.3

3. Enthalpy changes of reaction and formation

   Each named enthalpy change is one exact equation with its own definition. Reaction is the general case; formation builds one mole of a compound from its elements in their standard states.

   14 min 6.46.4(i)6.4(ii)

4. Enthalpy changes of combustion, neutralisation and atomisation

   Combustion counts per mole of fuel burnt, neutralisation per mole of water formed, atomisation per mole of gaseous atoms made. Know each definition and its equation, with state symbols.

   14 min 6.46.4(iii)6.4(iv)6.4(v)

5. Measuring the energy transferred

   A thermometer reading becomes an energy in joules through q = m c ΔT. The mass is always the mass of the water or solution that changed temperature, never the mass of the reactant.

   12 min 6.56.5(i)

6. From joules to kJ mol⁻¹

   Divide q by 1000, divide by the moles of the limiting reactant, then attach the sign from the direction of the temperature change. Worked for a solution in a cup and for a spirit burner.

   14 min 6.56.5(ii)

7. Errors, uncertainty and assumptions

   To evaluate a calorimetry result, name the source of error, say which way it moves the answer, and suggest an improvement. Uncertainty is the limit of your apparatus, written as a percentage; assumptions are what the calculation takes for granted.

   14 min 6.8

8. Graphs and cooling-curve corrections

   The highest thermometer reading is too low, because the mixture is already cooling. Plot temperature against time, extend the cooling line back to the moment of mixing, and read ΔT from that.

   13 min 6.8

9. Hess's Law and enthalpy cycles

   The enthalpy change of a reaction does not depend on the route. Draw a cycle with the route you want on one side and a route you have data for on the other, then add and subtract along the arrows.

   13 min 6.66.6(i)

10. Hess cycles from formation and combustion data

    With formation data, ΣΔfH(products) − ΣΔfH(reactants). With combustion data, ΣΔcH(reactants) − ΣΔcH(products). Draw the cycle first, check the arrows, then do the sum.

    15 min 6.66.6(ii)

11. Core Practical 2: method

    Magnesium cannot be burned in a calorimeter to give a clean enthalpy change of formation. Instead, measure the heat of magnesium with acid and of magnesium oxide with acid, then use Hess's Law.

    11 min 6.7

12. Core Practical 2: results, cycle and evaluation

    Turn two temperature changes into two enthalpy changes, close the cycle with the formation of water, and judge the answer against the data-book value and the uncertainties.

    15 min 6.76.66.6(ii)6.8

13. Bond enthalpies and enthalpy changes

    Breaking bonds takes in energy and making bonds gives it out. Add up the bonds broken, subtract the bonds made, and you have an estimate of ΔH for a reaction of gases.

    14 min 6.9

14. Limits of bond enthalpy calculations, and finding a mean bond enthalpy

    Bond enthalpies give a close estimate, not an exact answer, because the values are averages and the method only works for gases. Run the calculation backwards and a known ΔH gives a mean bond enthalpy.

    14 min 6.96.10

15. Which bond breaks first?

    A lower bond enthalpy means a weaker bond that needs less energy to break, so reactions that start by breaking it tend to go faster at room temperature. This is a guide, not a guarantee.

    9 min 6.11

Also in this topic: Topic 6 on a page

Spec checklist

Every point you can be examined on in this topic, and where it is explained. Tick them off as you go.