Understand exothermic and endothermic reactions, interpret energy diagrams, define enthalpy change, apply Hess’s Law, and relate bond energy to heat of reaction.
In an exothermic reaction, energy is released to the surroundings. The products have less stored chemical energy (lower enthalpy) than the reactants. The surroundings therefore gain energy and become warmer.
In an endothermic reaction, energy is absorbed from the surroundings. The products have more stored chemical energy than the reactants. The surroundings lose energy and become cooler.
Enthalpy (H) is the total heat energy stored in a substance at constant pressure. We cannot measure absolute enthalpy, but we can measure the change in enthalpy:
Enthalpy is measured in kJ·mol−1. The sign of ΔH tells us the direction of energy flow relative to the system.
A potential energy (PE) diagram (also called an energy profile) shows how the potential energy of the reacting system changes as the reaction proceeds. Key features:
| Feature | Exothermic | Endothermic |
|---|---|---|
| Reactant vs product energy | Reactants higher | Products higher |
| ΔH value | Negative (< 0) | Positive (> 0) |
| Energy gap | Reactants − products | Products − reactants |
| Activation energy (Ea) | Energy from reactants to peak | Energy from reactants to peak |
The activated complex (transition state) sits at the peak of the energy profile. It is an unstable, high-energy arrangement of atoms that exists momentarily between reactants and products.
A catalyst provides an alternative reaction pathway with a lower activation energy. This means more particles have sufficient energy to react, so the reaction is faster.
The standard enthalpy of formation is the enthalpy change when one mole of a compound is formed from its constituent elements in their standard states at 298 K and 101.3 kPa.
Hess’s Law states that the total enthalpy change for a reaction is independent of the pathway — it depends only on the initial reactants and final products. This allows us to calculate ΔH for reactions that cannot be measured directly.
Rules for manipulating thermochemical equations:
Energy is required to break bonds (endothermic process) and energy is released when bonds form (exothermic process). The heat of reaction can be estimated using average bond energies:
If more energy is released forming bonds than is needed to break bonds, ΔH < 0 (exothermic). Bond energy calculations give approximate answers because average values are used.
A calorimeter measures the heat released or absorbed by a chemical reaction by monitoring the temperature change of a known mass of water (or solution). The formula used is:
| Symbol | Quantity | Units |
|---|---|---|
| Q | Heat energy transferred | J or kJ |
| m | Mass of solution (usually taken as mass of water) | g |
| c | Specific heat capacity of water | 4.18 J·g−1·°C−1 |
| ΔT | Change in temperature (Tfinal − Tinitial) | °C or K |
To find ΔH per mole: calculate moles of limiting reactant, then ΔH = −Q / n (the negative sign because if the solution heats up, the reaction was exothermic).
The Born-Haber cycle is a Hess’s Law cycle used to calculate the lattice energy of ionic compounds. It connects several enthalpy changes:
By Hess’s Law: ΔH°f = sum of all steps in the cycle. Lattice energy is large and negative for ionic compounds, indicating strong electrostatic attraction between ions.
Entropy (S) is a measure of disorder or randomness in a system. Reactions tend toward higher entropy (ΔS > 0 means entropy increases).
Gibbs Free Energy (ΔG) combines enthalpy and entropy to predict spontaneity:
T is temperature in Kelvin. A reaction can be spontaneous at high temperatures if TΔS is large enough to overcome a positive ΔH.
| Time after mixing (min) | Temperature (°C) |
|---|---|
| 1 | 14.0 |
| 2 | 14.6 |
| 3 | 15.2 |
| 4 | 15.8 |
| 5 | 16.4 |