Exothermic Reaction ($\Delta H < 0$)

Products sit at LOWER energy level than reactants. Energy released to surroundings as heat.

$\text{Reactants} \xrightarrow{\Delta H < 0} \text{Products} + \text{Heat}$

Example: $CH_{4}$(g) + $2O_{2}$(g) → $CO_{2}$(g) + $2H_{2}O$(l), $\Delta H = -890$ kJ/mol

Endothermic Reaction ($\Delta H > 0$)

Products sit at HIGHER energy level. System absorbs heat from surroundings.

$\text{Reactants} + \text{Heat} \xrightarrow{\Delta H > 0} \text{Products}$

Example: $CaCO_{3}$(s) → CaO(s) + $CO_{2}$(g), $\Delta H = +178$ kJ/mol

Hess's Law Diagram

Hess's law states the total enthalpy change is path-independent:

$\Delta H_{total} = \Delta H_1 + \Delta H_2 + \Delta H_3 = \Delta H_{direct}$

even if the reaction proceeds through multiple intermediate steps.

Energy Profile: Exothermic vs Endothermic

Enthalpy (H)

Exothermic Reaction

Reactants ($H_{1}$) Products ($H_{2}$) Ea $\Delta H$ < 0 (negative)

Products at LOWER energy Heat released to surroundings

Endothermic Reaction

Reactants ($H_{1}$) Products ($H_{2}$) Ea $\Delta H$ > 0 (positive)

Products at HIGHER energy Heat absorbed from surroundings

Reaction Progress →

Sign Convention Summary

• Products lower than reactants → $\Delta H < 0$ (exothermic) → heat flows OUT
• Products higher than reactants → $\Delta H > 0$ (endothermic) → heat flows IN