Exothermic and Endothermic Reactions: MDCAT Chemistry notes
Exothermic and Endothermic Reactions MDCAT notes: sign of enthalpy change, heat content of reactants and products, bond energies and common examples.
Definitions
| Feature | Exothermic | Endothermic |
|---|---|---|
| Heat | Released to the surroundings | Absorbed from the surroundings |
| Sign of $\Delta H$ | Negative | Positive |
| Heat content | Reactants > products | Products > reactants |
| Temperature of surroundings | Rises | Falls |
| Energy profile | Products lie below reactants | Products lie above reactants |
$$\Delta H = H_{\text{products}} - H_{\text{reactants}}$$
If $H_{\text{products}} \lt H_{\text{reactants}}$, $\Delta H$ is negative (exothermic). If $H_{\text{products}} \gt H_{\text{reactants}}$, $\Delta H$ is positive (endothermic).
Why reactions release or absorb heat
Energy changes in chemical reactions come from bond breaking and bond formation:
- Breaking bonds absorbs energy (endothermic step).
- Forming bonds releases energy (exothermic step).
A reaction is exothermic when more energy is released in forming new bonds than is absorbed in breaking old ones. Combustion is exothermic for this reason. A reaction is endothermic when bond breaking needs more energy than bond formation gives out.
$$\Delta H \approx \sum(\text{bonds broken}) - \sum(\text{bonds formed})$$
Common examples
| Exothermic | Endothermic |
|---|---|
| Combustion of fuels, e.g. burning methane | $\mathrm{N_2 + O_2 \rightarrow 2NO}$ (formation of nitric oxide) |
| Neutralization of acid by alkali | Photosynthesis |
| Quicklime with water: $\mathrm{CaO + H_2O \rightarrow Ca(OH)_2}$ | Thermal decomposition, e.g. $\mathrm{CaCO_3 \rightarrow CaO + CO_2}$ |
| Respiration | Melting, evaporation, sublimation (e.g. dry ice) |
| Formation of most compounds from elements, e.g. $\mathrm{H_2O}$ | Dissolving $\mathrm{NH_4Cl}$ in water |
Nitrogen and oxygen in air
$\mathrm{N_2}$ and $\mathrm{O_2}$ exist together in the air without forming oxides of nitrogen because the reaction is strongly endothermic and the strong $\mathrm{N{\equiv}N}$ triple bond is hard to break. They combine only at very high temperatures, such as in lightning or inside car engines.
Exothermic reactions are more common
Most reactions that take place at ordinary temperatures are exothermic ($\Delta H$ negative). The products are at lower energy and therefore more stable than the reactants. Endothermic reactions usually need continuous heating to keep going.
Worked example
A reaction breaks bonds totalling $+1370\ \mathrm{kJ}$ and forms bonds totalling $-1850\ \mathrm{kJ}$. Then $\Delta H = 1370 - 1850 = -480\ \mathrm{kJ}$: the reaction is exothermic and the temperature of the surroundings rises.
Key formulas
- $\Delta H = H_{\text{products}} - H_{\text{reactants}}$
- $\Delta H \approx$ energy absorbed breaking bonds − energy released forming bonds
Common MDCAT traps
- In an exothermic reaction the reactants have more heat content than the products; in an endothermic reaction the products have more.
- Combustion is exothermic because more energy is released in bond making, not bond breaking.
- Endothermic reactions have a positive $\Delta H$.
- Sublimation of dry ice and formation of NO are endothermic; burning methane is exothermic.
- A rise in the temperature of the system (as with quicklime and water) signals an exothermic reaction.
Quick revision
- Exothermic: heat released, $\Delta H \lt 0$.
- Endothermic: heat absorbed, $\Delta H \gt 0$.
- Bond breaking absorbs energy; bond making releases it.
- $\mathrm{N_2}$ and $\mathrm{O_2}$ do not react in air because the reaction is endothermic.
- Slaking of lime is exothermic.