3.1Physical Chemistry
Energetics — Edexcel Linear - OL (Chemistry) Revision Notes & Topical Questions
5 min read 14 parts Lesson 4 of 9 Unit quiz
Energetics: Exothermic, Endothermic and Energy Profiles — EDEXCEL LINEAR – OL (CHEMISTRY) notes: Describe energy transfers in chemical reactions, draw energy profiles and carry out simple calorimetry practicals.Chemical reactions involve energy changes because bonds break and form.
- activation energy
- Worked example 1
- Worked example 2
- Worked example 3
- Worked example 4
- Calorimeter
Section 1 of 14Energy in chemical reactions
Chemical reactions involve energy changes because bonds break and form. The total energy of a closed system is conserved, so energy lost by the reacting chemicals is gained by the surroundings. Reactions are classified as exothermic (energy to the surroundings, temperature rises) or endothermic (energy from the surroundings, temperature falls).
| Exothermic examples | Endothermic examples |
|---|---|
| Burning fuels, respiration | Photosynthesis (light energy absorbed) |
| Neutralisation of acid with alkali | Thermal decomposition of limestone |
| Reaction of metals with acid | Dissolving ammonium nitrate or potassium chloride |
| Rusting (slow), hand warmers | Reaction of citric acid with sodium hydrogencarbonate |
Section 2 of 14Energy profile diagrams
An energy profile compares the energy of reactants and products.
Section 3 of 14Practical: measuring temperature changes
Polystyrene is a good insulator and the lid reduces heat loss, giving more accurate results. For a fuel, burn it in a spirit burner under a copper can containing a known mass of water and measure the temperature rise. Results lower than expected arise from heat loss, incomplete combustion and evaporation.
Section 4 of 14Comparing fuels
To compare fuels fairly
- use the same mass of water and the same distance between the flame and can, burn the same mass of fuel, and calculate energy per gram.
- Alcohol fuels such as ethanol give less energy per gram than hydrocarbons like octane but are renewable if made from sugar cane;
- hydrogen releases the most energy per gram and produces only water.
Section 5 of 14Energy changes in everyday life and industry
Energy changes matter in daily life.
- while hydrogen releases only water but is difficult to store.
The idea of conservation of energy underlies all of this: energy is not created or destroyed, only transferred from the chemical store of the reactants to thermal energy of the surroundings (exothermic), or from the surroundings to the chemical store of the products (endothermic).
Section 6 of 14Worked examples
Sodium hydroxide solution is added to hydrochloric acid and the temperature rises from 20 °C to 27 °C. Classify the reaction and explain. [2]2 marks
Exothermic: energy is released to the surroundings, so the temperature rises.
Draw the energy profile for an endothermic reaction. [3]3 marks
Products higher than reactants; a hump above both labelled activation energy; an upward arrow marked ΔH positive.
A student wants to improve accuracy in the polystyrene cup experiment. Suggest two changes. [2]2 marks
Use a lid and a double cup or insulating jacket, and stir to get an even temperature.
Give one use of an endothermic reaction. [1]1 mark
Instant cold packs using dissolving ammonium nitrate.
Section 7 of 14Bond breaking and bond making
Chemical reactions involve rearranging atoms. Energy is needed to break bonds in the reactants (endothermic) and energy is released when new bonds form in the products (exothermic). The overall energy change depends on which is larger:
| Condition | Overall | ΔH |
|---|---|---|
| Energy released making bonds > energy needed breaking bonds | Exothermic | Negative |
| Energy needed breaking bonds > energy released making bonds | Endothermic | Positive |
The bond energy is the energy needed to break one mole of a particular covalent bond in gaseous molecules, in kJ/mol. Stronger bonds have higher bond energies, such as C≡C and N≡N.
Section 8 of 14Step-by-step method
Example
- N2 + 3H2 → 2NH3 with N≡N 945, H–H 436, N–H 391 kJ/mol.
- Broken: 945 + 3 × 436 = 2253.
- Formed: 6 × 391 = 2346.
- ΔH = 2253 − 2346 = −93 kJ/mol, exothermic.
- Check: 2 mol of NH3 are formed, so for one mole it is −46.5 kJ/mol.
Section 9 of 14Energy level diagram link and uses
On an energy profile the difference between reactants and products represents ΔH, which can be calculated from bond energies.
- and predicting whether a reaction would be hazardous.
- (such as explosives that release a huge amount of energy as strong bonds in nitrogen gas form)
Section 10 of 14Comparing bond energies and drawing the energy cycle
An alternative layout for the calculation is an energy cycle: reactants to separate atoms (energy in, bonds broken), then separated atoms to products (energy out, bonds formed). ΔH is the sum of the two steps with the correct signs. Many students prefer a table.
| Bond | Number | Energy each (kJ/mol) | Total (kJ/mol) |
|---|---|---|---|
| H–H (broken) | 3 | 436 | 1308 |
| N≡N (broken) | 1 | 945 | 945 |
| N–H (formed) | 6 | 391 | 2346 |
Total broken 2253, total formed 2346, ΔH = −93 kJ/mol.
- which explains the use of high temperature and a catalyst in the Haber process even though the overall change is exothermic.
- so the experimental ΔH is more negative.
Section 11 of 14Worked examples
Calculate ΔH for 2H2 + O2 → 2H2O using H–H 436, O=O 498, O–H 464. [4]4 marks
Broken: 2 × 436 + 498 = 1370. Formed: 4 × 464 = 1856. ΔH = 1370 − 1856 = −486 kJ/mol.
Explain why the combustion of fuels is exothermic in terms of bonds. [2]2 marks
More energy is released forming the strong C=O and O–H bonds than is needed to break the bonds in the fuel and oxygen.
A reaction breaks bonds requiring 1500 kJ and forms bonds releasing 1200 kJ. Calculate ΔH and classify. [2]2 marks
ΔH = 1500 − 1200 = +300 kJ/mol; endothermic.
Why are bond energies only an approximation? [1]1 mark
They are averages over many molecules, not exact for a specific compound.
Section 12 of 14Key terms
| Term | Meaning |
|---|---|
| Exothermic | Releases energy to the surroundings |
| Endothermic | Absorbs energy from the surroundings |
| Activation energy | Energy needed to start the reaction |
| Energy profile | Diagram of energy changes during a reaction |
| Calorimeter | Insulated container for measuring heat changes |
| Surroundings | Everything outside the reacting chemicals |
| Term | Meaning |
|---|---|
| Bond energy | Energy to break one mole of a bond |
| Enthalpy change | ΔH, energy change of a reaction |
| Bonds broken | Reactant bonds, energy absorbed |
| Bonds formed | Product bonds, energy released |
| Displayed formula | Formula showing all bonds |
| Exothermic | ΔH negative |
Section 13 of 14Syllabus checklist
CH4 + 2O2 → CO2 + 2H2O.
- Bonds formed
- 2 C=O (2 × 805 = 1610) and 4 O–H (4 × 464 = 1856), total 3466.
- ΔH = 2648 − 3466 = −818 kJ/mol.
Section 14 of 14Exam technique
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