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IGCSE Chemistry 0620 · Chapter 5 notes · ~9 min read

Chemical energetics

Why reactions get hot or cold, how to read an energy-profile diagram, and how to calculate the energy change of a reaction from bond energies.

Exothermic & endothermicEnergy profilesΔHActivation energyBond energies
RSby Riyaz Sir
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No tag = Core (everyone)EXTENDED Extended onlyhighlight = comes up again and againTest yourself at the end

In 30 seconds

  • Exothermic reactions give energy out to the surroundings — the temperature rises. Endothermic reactions take energy in — the temperature falls.
  • An energy profile shows the energy of reactants and products. The gap between them is ΔH. EXT
  • Activation energy (Ea) is the energy hill reactants must climb before reacting. EXT
  • EXT ΔH = energy to break bonds − energy released making bonds. Breaking takes in energy. Making releases it.

By the end you can

  • Define exothermic and endothermic reactions and give examples.
  • Read an energy-profile diagram and say whether a reaction is exothermic or endothermic.
  • Label activation energy and ΔH on a profile and give the sign of ΔH. EXT
  • Calculate ΔH from bond energies. EXT

1Exothermic and endothermic reactions

Heat out or heat in

Exothermic reaction — a reaction that transfers thermal energy to the surroundings. The temperature of the surroundings rises.

Endothermic reaction — a reaction that takes in thermal energy from the surroundings. The temperature of the surroundings falls.

ExothermicEndothermic
energygiven outtaken in
temperature of surroundingsgoes upgoes down
products compared with reactantslower in energyhigher in energy
ΔH EXTnegative (−)positive (+)
examplesburning fuels, neutralisation, respiration, rusting, adding water to quicklimethermal decomposition (CaCO₃ → CaO + CO₂), photosynthesis, dissolving ammonium nitrate in water, citric acid + sodium hydrogencarbonate
Memory trickEXO = EXIT — energy exits the reaction. ENDO = energy goes IN.
Picture itA campfire is exothermic: you feel heat leaving it. A cold pack is endothermic: it pulls heat out of your skin, so it feels cold.

How to measure it

  • Put the reactants in a polystyrene cup (an insulator, so little heat is lost) with a lid, stir, and use a thermometer to measure the temperature at the start and the highest (or lowest) temperature.
  • Rise in temperature → exothermic. Fall in temperature → endothermic.
Exam tipAlways describe the change in the temperature of the surroundings (or the mixture), not "the reaction gets hot". Give numbers if you are asked to describe results: "temperature rose from 21 °C to 27 °C".
Common mistake"Endothermic reactions get hot because they need heating." They need heat from the surroundings, so the temperature of the surroundings falls.

2Energy profile diagrams

Reading a reaction pathway

EXOTHERMICENDOTHERMICprogress of reactionenergyprogress of reactionenergyreactantsproductsreactantsproductsΔHEaΔHEaproducts have less energyenergy released to the surroundingsΔH is negative (−)products have more energyenergy taken in from the surroundingsΔH is positive (+)
Energy profiles. Left: exothermic — products are lower than reactants. Right: endothermic — products are higher.
  • The vertical axis is energy. The horizontal axis is progress of reaction (from reactants to products).
  • Exothermic: the products are lower than the reactants. The difference is energy released to the surroundings.
  • Endothermic: the products are higher than the reactants. The difference is energy taken in from the surroundings.
  • EXT ΔH is drawn as a vertical arrow between the reactants level and the products level.
  • EXT Activation energy (Ea) is the vertical arrow from the reactants up to the top of the hump.

Activation energy (Ea) EXT — the minimum energy that colliding particles need in order to react.

EXT ΔH = energy of products − energy of reactants
negative for exothermic, positive for endothermic
  • EXT Both exothermic and endothermic reactions have an activation energy. The hump is the energy needed to start breaking bonds.
Picture itA boulder in a dip on a hillside: you must push it up over the hump (Ea) before it can roll down to the valley (an exothermic reaction). Even when the valley is much lower, the push is still needed — a match is needed to light a candle.
  • EXT A catalyst lowers the hump (it gives an alternative route with lower Ea), but ΔH stays the same. See the rates chapter.
Worked example

EXT A reaction has reactants at 200 kJ, a peak at 350 kJ and products at 120 kJ. Find Ea and ΔH, and say whether it is exothermic.

  1. Ea = peak − reactants = 350 − 200 = 150 kJ
  2. ΔH = products − reactants = 120 − 200 = −80 kJ
  3. ΔH is negative and the products are lower, so it is exothermic
Exam tipIn a "draw a profile" question: (1) label the axes "energy" and "progress of reaction"; (2) draw reactants and products at the right relative levels; (3) draw the hump; (4) show Ea from the reactants and ΔH between the levels, each with an arrow.
Common mistakeDrawing Ea from the products or measuring it down to zero. Ea starts at the reactants level.
Common mistakeWriting a positive ΔH for an exothermic reaction. Exothermic = negative ΔH.

3Bond energies EXT

Breaking takes energy, making releases it

Bond energy — the energy needed to break one mole of a particular covalent bond, in kJ/mol.

  • Breaking bonds is endothermic — energy must be put in to pull atoms apart.
  • Making bonds is exothermic — energy is released when atoms join.
  • The overall energy change depends on which is bigger:
  • more energy released making new bonds than taken in breaking old ones → exothermic
  • more energy taken in breaking bonds than released making new ones → endothermic
ΔH = Σ (energy of bonds broken) − Σ (energy of bonds made)
in kJ/mol
H₂ + Cl₂2H + 2Cl (separate atoms)2 HClbreak bonds436 + 242= +678 kJmake bonds2 × 431= −862 kJΔHΔH = energy to break bonds − energy released making bonds= 678 − 862 = −184 kJ/mol → exothermicbond energies: H–H 436 · Cl–Cl 242 · H–Cl 431 kJ/mol
Hydrogen + chlorine → hydrogen chloride. Breaking the bonds costs 678 kJ. Making new bonds gives back 862 kJ, so the reaction is exothermic.
Worked example

EXT Calculate ΔH for H₂ + Cl₂ → 2HCl. Bond energies: H–H = 436, Cl–Cl = 242, H–Cl = 431 kJ/mol.

  1. Bonds broken: 1 × H–H + 1 × Cl–Cl = 436 + 242 = 678 kJ
  2. Bonds made: 2 × H–Cl = 2 × 431 = 862 kJ
  3. ΔH = 678 − 862 = −184 kJ/mol
  4. Negative, so it is exothermic
Worked example

EXT Calculate ΔH for CH₄ + 2O₂ → CO₂ + 2H₂O. Bond energies: C–H = 413, O=O = 498, C=O = 805, O–H = 464 kJ/mol.

  1. Bonds broken: 4 × C–H = 4 × 413 = 1652; 2 × O=O = 2 × 498 = 996
  2. Total broken = 1652 + 996 = 2648 kJ
  3. Bonds made: 2 × C=O = 2 × 805 = 1610; 4 × O–H = 4 × 464 = 1856
  4. Total made = 1610 + 1856 = 3466 kJ
  5. ΔH = 2648 − 3466 = −818 kJ/mol (exothermic)
Worked example

EXT Calculate ΔH for N₂ + O₂ → 2NO. Bond energies: N≡N = 945, O=O = 498, N=O (in NO) = 631 kJ/mol.

  1. Bonds broken: 945 + 498 = 1443 kJ
  2. Bonds made: 2 × 631 = 1262 kJ
  3. ΔH = 1443 − 1262 = +181 kJ/mol
  4. Positive, so it is endothermic — this is why nitrogen and oxygen only react at the very high temperatures inside a car engine
Exam tipCount the bonds for every molecule: draw the displayed formula. In CH₄ there are 4 C–H bonds; in H₂O there are 2 O–H bonds; in CO₂ there are 2 C=O bonds. Then do broken − made.
Common mistakeSubtracting the wrong way round. If your answer for a burning fuel is positive, check again — fuels release energy, so ΔH should be negative.
Common mistakeSaying that breaking bonds releases energy. It does not — breaking always takes energy in.

4Test yourself

Test yourself — cover the answers first

Quick check

  1. Give two examples of exothermic reactions.
  2. What happens to the temperature of the surroundings in an endothermic reaction?
  3. Name one endothermic process.
  4. Are the products higher or lower in energy than the reactants in an exothermic reaction?
  5. EXT What is the sign of ΔH for an endothermic reaction?
  6. EXT Define activation energy.
  7. EXT Is breaking a bond exothermic or endothermic?

Exam-style

  1. EXT Draw and label an energy profile for an exothermic reaction. [4]
  2. EXT Calculate ΔH for 2H₂ + O₂ → 2H₂O. Bond energies: H–H 436, O=O 498, O–H 464 kJ/mol. [3]
  3. EXT A student says "Exothermic reactions do not need any energy to start". Explain why this is wrong. [2]
  4. Describe an experiment to show that a reaction is exothermic. [3]
Show answers
  1. any two from: burning, neutralisation, respiration, rusting
  2. it falls
  3. thermal decomposition of calcium carbonate / photosynthesis / dissolving ammonium nitrate
  4. lower
  5. positive (+)
  6. the minimum energy colliding particles must have in order to react
  7. endothermic
  8. axes: energy (vertical) and progress of reaction (horizontal)
    reactants level higher than products level
    hump between them
    Ea arrow from reactants to the top of the hump, and ΔH arrow (downwards) from reactants to products
  9. broken: 2 × 436 + 498 = 872 + 498 = 1370 kJ
    made: 4 × 464 = 1856 kJ
    ΔH = 1370 − 1856 = −486 kJ/mol
  10. bonds in the reactants must be broken first, which takes energy
    this is the activation energy, e.g. a match is needed to light a fuel
  11. mix the reactants in a polystyrene (insulated) cup
    measure the temperature with a thermometer before and after, stirring
    the temperature goes up
Practise this chapter
Exothermic & endothermic reactions
Next chapterRates, equilibrium & redox →
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