Enthalpy change and reaction pathways
Enthalpy change, ΔH
- Every reaction takes in or gives out energy.
- That energy change (at constant pressure) is the enthalpy change, $\Delta H$.
- Its sign tells you the direction of energy flow.
Exothermic vs endothermic
- Exothermic: gives out heat → products lower than reactants → $\Delta H < 0$.
- Endothermic: takes in heat → products higher → $\Delta H > 0$.
- The "hill" between them is the activation energy — the least energy needed to react.
Practice
In an exothermic reaction:
Exothermic reactions release heat; products sit below reactants, so ΔH is negative.
Practice
The activation energy is:
Activation energy is the height of the "hill" — the minimum energy needed to start the reaction.
Standard conditions and types
- Values are compared under standard conditions ($298\ \text{K}$, $101\ \text{kPa}$), shown by $^{\ominus}$.
| Symbol | Enthalpy change of |
|---|---|
| $\Delta H_f^{\ominus}$ | formation — 1 mole of compound from its elements |
| $\Delta H_c^{\ominus}$ | combustion — 1 mole burns completely in oxygen |
| $\Delta H_{\text{neut}}^{\ominus}$ | neutralisation — 1 mole of water from acid + alkali |
Practice
The standard enthalpy change of combustion is for:
ΔHc is per mole of substance burned completely; ΔHf is formation from elements; ΔH_neut makes one mole of water.
Practice
Standard conditions for enthalpy values are:
Standard conditions are 298 K and 101 kPa, with each substance in its normal state (symbol ⊖).
You've got it
Key idea
- exothermic $\Delta H < 0$ (products lower); endothermic $\Delta H > 0$ (products higher)
- activation energy = the energy hill, the least needed to react
- standard conditions: 298 K, 101 kPa ($^{\ominus}$)
- types: formation ($\Delta H_f$), combustion ($\Delta H_c$), neutralisation ($\Delta H_{\text{neut}}$)