Exothermic & Endothermic Reactions

Energy is conserved in chemical reactions. The amount of energy in the universe at the end of a chemical reaction is the same as before the reaction takes place. If a reaction transfers energy to the surroundings the product molecules must have less energy than the reactants, by the amount transferred.

Exothermic

Energy released to surroundings → temperature increases.

• Examples: combustion, neutralisation, oxidation reactions.

• Reaction profile: products lower in energy than reactants.

• ΔH is negative (because energy leaves the system).

Endothermic

• Energy taken in from surroundings → temperature decreases.

• Examples: thermal decomposition, citric acid + sodium hydrogen carbonate, photosynthesis.

• Reaction profile: products higher in energy.

• ΔH is positive.

Grade 9 insight:

If you release more energy (from making bonds) than you take in (breaking bonds) → EXOTHERMIC.

If you take in more energy (breaking bonds) than you release (making bonds) → ENDOTHERMIC.

Breaking bonds = endothermic

(energy must be put in)

Making bonds = exothermic

(energy released)

Reaction Profiles (Energy Diagrams)

• Reactants

• Products

• Activation energy (Ea) → minimum energy needed to start the reaction.

• Overall energy change (ΔH) → difference between reactants and products.

  

Endothermic → absorbs energy → solution loses heat → colder
Exothermic → releases energy → solution gains heat → hotter

Bond Energy Calculations

Energy change =
Energy required to break bonds – Energy released when bonds form

Because:

• Breaking bonds = endothermic (requires energy)

• Making bonds = exothermic (releases energy)

Example (Grade 9 standard):

H₂ + Cl₂ → 2HCl
Bond energies:
H–H = 436 kJ/mol
Cl–Cl = 243 kJ/mol
H–Cl = 431 kJ/mol

Calculate ΔH:

1. Energy in (breaking):
   H–H + Cl–Cl = 436 + 243 = 679 kJ

2. Energy out (forming):
   2 × H–Cl = 2 × 431 = 862 kJ

3. ΔH = 679 – 862 = –183 kJ/mol (exothermic)

What’s happening physically

• Chemical reactions involve bond breaking and bond forming:

  • Breaking bonds = energy absorbed → endothermic

  • Forming bonds = energy released → exothermic

• In a solution:

  • Energy absorbed from water → solution cools → endothermic

  • Energy released into water → solution warms → exothermic

💡 Key idea: the solution acts as the surroundings, so its temperature change tells you the energy flow.

Experimental setup

• Use a polystyrene cup → minimizes energy loss to air

• Thermometer → measures solution temperature precisely

• Keep variables controlled:

  • Volume of reactants

  • Concentration of solutions

  • Initial temperature

  • Surface area of solids (if used)

• Measure temperature at regular intervals → can plot temperature vs time graph for accuracy

How to tell endo vs exo

• Endothermic: temperature drops → reaction absorbs energy from the solution

• Exothermic: temperature rises → reaction releases energy into the solution

Grade 9 phrasing:

• Always reference energy transfer; saying “it feels colder” is not enough for full marks.

Data interpretation

• Use temperature change (ΔT) to calculate approximate energy change in kJ if given volume and specific heat capacity.

• Graph tip: steep drop/rise = fast reaction, gradual = slower reaction.