Exothermic vs endothermic reactions — examples and energy diagrams

Question

Distinguish between exothermic and endothermic reactions. Give two examples of each, and describe what their energy diagrams (reaction profiles) look like.

Solution — Step by Step

An exothermic reaction is one that releases energy (usually as heat) to the surroundings. The products have lower energy than the reactants. The system loses energy → the surroundings gain energy → the surroundings get warmer.

The enthalpy change $\Delta H$ for an exothermic reaction is negative ( $\Delta H < 0$ ).

1. Combustion of natural gas (methane):

\text{CH}_4(g) + 2\text{O}_2(g) \rightarrow \text{CO}_2(g) + 2\text{H}_2\text{O}(l) \quad \Delta H = -890 \text{ kJ/mol}

The hot flame you see on a gas stove is the released energy. The surroundings (air, pot, food) get heated.

2. Neutralisation of a strong acid with a strong base:

\text{HCl}(aq) + \text{NaOH}(aq) \rightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l) \quad \Delta H \approx -57 \text{ kJ/mol}

Mix these in a flask — the solution noticeably warms up. The reaction releases heat.

3. Respiration (cellular): Glucose oxidation releases ~2870 kJ/mol — this energy powers all life processes.

An endothermic reaction is one that absorbs energy from the surroundings. The products have higher energy than the reactants. The surroundings cool down as energy flows into the reaction.

The enthalpy change $\Delta H$ for an endothermic reaction is positive ( $\Delta H > 0$ ).

1. Thermal decomposition of calcium carbonate (limestone):

\text{CaCO}_3(s) \xrightarrow{\Delta} \text{CaO}(s) + \text{CO}_2(g) \quad \Delta H = +178 \text{ kJ/mol}

This reaction requires continuous heating in a kiln to proceed. The furnace must constantly supply energy — it doesn’t self-sustain.

2. Dissolution of ammonium nitrate in water:

\text{NH}_4\text{NO}_3(s) \xrightarrow{H_2O} \text{NH}_4^+(aq) + \text{NO}_3^-(aq) \quad \Delta H = +25 \text{ kJ/mol}

The water gets cold — this is the principle behind instant cold packs used in sports injuries. Energy is absorbed from the water to break ionic bonds in NH₄NO₃.

3. Photosynthesis: Plants absorb light energy to convert CO₂ and H₂O into glucose — strongly endothermic ( $\Delta H = +2870$ kJ/mol).

An energy profile diagram plots potential energy (y-axis) against reaction progress (x-axis).

Exothermic reaction profile:

Reactants at a higher energy level
Products at a lower energy level
The “hill” in between is the activation energy (energy needed to start the reaction)
The products are DOWN the energy slope from reactants
$\Delta H$ = negative (shown as a drop from reactants to products)

Endothermic reaction profile:

Reactants at a lower energy level
Products at a higher energy level
Still has an activation energy hill
The products are UP the energy slope from reactants
$\Delta H$ = positive (shown as a rise from reactants to products)

Why This Works

The energy change in a reaction comes from bond breaking and bond forming:

Breaking bonds requires energy (endothermic step)
Forming bonds releases energy (exothermic step)

If more energy is released in forming new bonds (products) than is absorbed in breaking old bonds (reactants) → net exothermic ( $\Delta H < 0$ ). If more energy is absorbed in breaking old bonds than is released in forming new bonds → net endothermic ( $\Delta H > 0$ ).

Every chemical reaction involves both bond breaking and bond forming. The sign of $\Delta H$ tells you which won.

Alternative Method — Temperature Test

A quick practical test:

Mix chemicals in an insulated container
If temperature rises: exothermic (heat released to surroundings = container)
If temperature falls: endothermic (heat absorbed from surroundings = container cools)

This is how calorimetry works — measure temperature change to determine $\Delta H$ .

Common Mistake

Students sometimes confuse the sign of $\Delta H$ with whether the reaction “needs heat to start.” ALL reactions need some activation energy to start — even exothermic ones. Burning methane (exothermic) still needs a spark to start. The sign of $\Delta H$ tells you whether heat is released or absorbed overall, not whether a trigger is needed.

Also, do not confuse “endothermic” with “endergonic” — these are different thermodynamic terms. Endothermic refers to enthalpy (heat content); endergonic refers to Gibbs free energy. An endothermic reaction can be spontaneous if entropy increases enough. These distinctions appear at Class 11 level.

Memory trick: “Exo” = exit, energy exits the system (released). “Endo” = inside, energy enters the system (absorbed). The system is the reaction; the surroundings are everything else (the flask, the air, you).

Question

Solution — Step by Step

Why This Works

Alternative Method — Temperature Test

Common Mistake

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