Exothermic and endothermic reactions

Enthalpy

• Enthalpy (H) is the chemical energy stored in a substance.
• You cannot measure enthalpy itself, but you can measure the change in it during a reaction.
• $\Delta_rH$ is the enthalpy change of reaction — the energy taken in or given out.
• Its unit is $kJ\ mol^{-1}$.

The two kinds of reaction

	Exothermic	Endothermic
Energy is	released to the surroundings	absorbed from the surroundings
Temperature of surroundings	rises — feels hot	falls — feels cold
Products compared with reactants	lower enthalpy	higher enthalpy
Sign of $\Delta_rH$	negative (−)	positive (+)

• The sign follows from where the energy went. In an exothermic reaction the products have less stored energy than the reactants, so the change is a decrease — negative.

Reading an enthalpy diagram

• The vertical axis is enthalpy; reactants are on the left and products on the right.
• $\Delta_rH$ is the vertical gap between the reactant level and the product level.
• Products lower than reactants → exothermic, $\Delta_rH$ negative.
• Products higher than reactants → endothermic, $\Delta_rH$ positive.

Everyday examples

Exothermic:

• Combustion — burning any fuel. Always exothermic.
• Neutralisation of an acid with an alkali.
• Respiration.
• Most reactions of metals with acids.

Endothermic:

• Thermal decomposition — heating limestone to make lime, for instance.
• Photosynthesis.
• Dissolving ammonium nitrate in water — this is how instant cold packs work.

Writing $\Delta_rH$ properly

• Always give the sign. A missing minus sign turns an exothermic reaction into an endothermic one.
• Always give the unit, $kJ\ mol^{-1}$.
• Write it beside the equation:

$$CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O \qquad \Delta_rH = -890\ kJ\ mol^{-1}$$

• If the equation is reversed, the sign of $\Delta_rH$ flips but the number stays the same.

The wording trap

• The reaction releases energy; the surroundings get hotter.
• So an exothermic reaction feels hot, but its $\Delta_rH$ is negative — because it is the chemicals that have lost energy.
• Students often reason "it got hotter, so energy went up, so positive". That is backwards: $\Delta_rH$ describes the chemicals, not the beaker.