Endothermic and exothermic reactions Notes

Endothermic vs Exothermic – What Happens to Energy and Temperature?

Endothermic Reactions

1. The system (chemicals reacting) gains energy.
2. The surroundings lose energy, so they cool down.
3. In terms of bonding:
   1. Overall, more energy is required to break bonds in the reactants than is released when new bonds form in the products.
   2. The “extra” energy needed comes from the surroundings → energy is absorbed.
4. Temperature change of the surroundings
   1. The temperature of the surroundings decreases.
   2. If you touch the container, it feels colder.

Exothermic Reactions

1. The system loses energy,
2. The surroundings gain energy, so they warm up.
3. In terms of bonding:
   1. Overall, more energy is released when new bonds form in the products than is needed to break bonds in the reactants.
   2. The excess energy is transferred to the surroundings.
4. Temperature change of the surroundings
   1. The temperature of the surroundings increases.
   2. If you touch the container, it feels warmer.

Summary Table – Endothermic vs Exothermic

Energy Profile Diagrams – Showing Energy Changes on a Graph

Energy profile diagrams show how the energy of a system changes as a reaction proceeds from reactants to products.

Key features:

1. Reactant energy level
2. Product energy level
3. Activation energy (Eₐ) – the energy needed to start the reaction
4. Enthalpy change (ΔH) – the overall energy change between reactants and products

Enthalpy and Enthalpy Change

Enthalpy change (ΔH):

$$\Delta H=H_{\text {products }}-H_{\text {reactants }}$$

1. $ΔH < 0$ → exothermic (energy released).
2. $ΔH > 0$→ endothermic (energy absorbed).

Energy Profile Diagram – Exothermic Reaction

1. Features:
   1. Reactants start at a higher energy than products.
   2. The curve rises to a peak → this is the activation energy, $E_a$.
   3. The curve then falls to the product energy level.
2. $\Delta H$ (enthalpy change):
   1. The vertical difference between the reactant and product levels.
   2. For an exothermic reaction, products are lower than reactants → $\Delta H$ is negative.

Energy Profile Diagram – Endothermic Reaction

1. Features:
   1. Reactants start at a lower energy than products.
   2. The curve rises to a peak (activation energy, $E_a$).
   3. The products end at a higher energy than the reactants.
2. $\Delta H$ (enthalpy change):
   1. The vertical difference between reactants and products.
   2. For an endothermic reaction, products are higher than reactants → $\Delta H$ is positive.

Why Does Enthalpy Change Occur?

1. In any reaction:
   1. Bonds in reactants must be broken → this requires energy.
   2. New bonds form in products → this releases energy.
2. The enthalpy change depends on the balance between:
   1. Energy absorbed to break bonds.
   2. Energy released when new bonds form.
   3. If more energy is released than absorbed → reaction is exothermic ($\Delta H < 0$).
   4. If more energy is absorbed than released → reaction is endothermic ($\Delta H > 0$).

Everyday Examples – Where Do We See These Energy Changes?

Everyday Exothermic Processes

1. Combustion of fuels (gas, petrol, candles):
   1. Release heat and light → used for heating, cooking, engines.
2. Respiration in cells: $$\text { glucose }+ \text { oxygen } \text { → } \text { carbon dioxide }+ \text { water }+ \text { energy }$$
   1. Releases energy your body uses for movement, growth, and temperature control.
3. Hand warmers / self-heating cans:
   1. Use exothermic reactions (e.g. iron oxidation or CaO + H₂O) to release heat.

Everyday Endothermic Processes

1. Photosynthesis (plants absorbing sunlight to make glucose).
2. Instant cold packs used in sports injuries.
3. Dissolving some salts in water (e.g. ammonium nitrate) where the solution temperature drops.
4. Evaporation of sweat from your skin – absorbs heat, helping cool the body.