The standard enthalpy of formation, $\Delta H\degree_f$ , is the enthalpy change when one mole of a substance is formed from its elements in their standard states.

→ The process of making bonds releases energy (as atoms are going from a higher energy, more unstable state to a lower energy more stable state) and is an endothermic process. The process of breaking bonds takes energy (as atoms are going from a lower energy to a higher energy state) and is an exothermic process.

The standard state of a substance is its most stable state at a pressure of $1$ atmosphere and at a specified temperature, usually taken as $298K$.

The standard enthalpy of a reaction can be calculated from the standard enthalpies of formation of the reactants and products, as shown in the following equation. The units for enthalpy in this equation are $kJ \space mol^{-1}$.

$$ \Delta H\degree_f=\Sigma\Delta H\degree_{f(products)} \space - \space \Sigma\Delta H\degree_{f(reactants)} $$

The entropy ($S$) of a system is a measure of the degree of disorder of the system.

The change in standard entropy for a reaction system can be calculated from the standard entropies of the reactants and products, as shown in the following reaction. The *units for entropy in this equation are $kJ \space K^{-1} \space mol^{-1}$.

$$ \Delta S \degree = \Sigma S \degree _{(products)} \space - \space \Sigma S \degree _{(reactants)} $$

→ The standard entropy of a substance is the entropy value for the substance in its standard state.

The greater the degree of disorder, the greater the entropy.

Entropy increases as temperature increases.

→ Solids have low disorder and gases have high disorder.

→ There is a rapid increase in entropy at the melting point of a substance and an even more rapid and larger change in entropy at the boiling point.

The second law of thermodynamics states that the total entropy of a reaction system and its surroundings always increases for a spontaneous process.

Heat energy released by the reaction system into the surroundings increases the entropy of the surroundings — exothermic systems increase the entropy of the surroundings.

Heat energy absorbed by the reaction system from the surroundings decreases the entropy of the surroundings — endothermic systems decrease the entropy of the surroundings.

The third law of thermodynamics states that the entropy of a perfect crystal at $0K$ is zero.

The change in free energy for a reaction is related to the enthalpy and entropy changes, as shown in the following equation.

$$ \Delta G \degree = \Delta H \degree - \space T \Delta S \degree $$

→ In this equation, $\Delta S \degree$ must be input in $kJ \space K^{-1} \space mol^{-1}$ and when $\Delta S \degree$ is calculated, the resulting number will also be in $kJ \space K^{-1} \space mol^{-1}$. However, questions will often ask the answer to be given in $J \space K^{-1} \space mol^{-1}$, in which case the previous result must be multiplied by $1000$.

If the change in free energy ($\Delta G \degree$) between reactants and products is negative, a reaction may occur and the reaction is said to be feasible. A feasible reaction is one that tends towards the products rather than the reactants. This does not give any indication of the rate of the reaction.