Enthalpy Change of a Reaction and Types of Enthalpy Change

The overall enthalpy of reaction, $\Delta_r H^\circ$, of any chemical equation can be broken down into a family of precisely defined, standardized enthalpy changes, each describing one very specific type of process carried out under standard conditions (usually $298\ \text{K}$ and $1\ \text{bar}$ pressure, denoted by the superscript $^\circ$). Learning these named quantities lets a chemist look up and combine tabulated data to calculate the enthalpy change of almost any reaction, even one that has never actually been carried out in a laboratory calorimeter, using Hess's law (covered in the next section).

• Standard enthalpy of formation ($\Delta_f H^\circ$) — the enthalpy change when one mole of a compound is formed from its constituent elements, each in their most stable standard state. By convention, $\Delta_f H^\circ$ of every element in its standard state is exactly zero. Example: $\text{C}(s, \text{graphite}) + \text{O}_2(g) \rightarrow \text{CO}_2(g)$, $\Delta_f H^\circ = -393.5\ \text{kJ mol}^{-1}$.
• Standard enthalpy of combustion ($\Delta_c H^\circ$) — the enthalpy change when one mole of a substance is completely burnt (oxidized) in excess oxygen. Combustion enthalpies are always negative (exothermic) and are readily measured using a bomb calorimeter. Example: $\text{CH}_4(g) + 2\text{O}_2(g) \rightarrow \text{CO}_2(g) + 2\text{H}_2\text{O}(l)$, $\Delta_c H^\circ = -890.3\ \text{kJ mol}^{-1}$.
• Standard enthalpy of atomization ($\Delta_a H^\circ$) — the enthalpy required to convert one mole of a substance completely into gaseous atoms. It is always positive (endothermic), since breaking all bonds requires energy input. Example: $\text{CH}_4(g) \rightarrow \text{C}(g) + 4\text{H}(g)$, $\Delta_a H^\circ = +1665\ \text{kJ mol}^{-1}$.
• Bond dissociation enthalpy — the energy required to break one mole of a specific covalent bond in a gaseous molecule, producing gaseous fragments; dividing an atomization enthalpy by the number of identical bonds broken gives the average bond enthalpy for that bond type.
• Enthalpy of ionization — the energy required to remove one mole of electrons from one mole of gaseous atoms (first ionization enthalpy) or from one mole of already-singly-charged gaseous cations (second ionization enthalpy, always larger, since removing an electron from an already-positive ion is harder). Both are always positive.
• Enthalpy of solution ($\Delta_{sol} H^\circ$) — the enthalpy change when one mole of a solute is dissolved in a large (effectively infinite) quantity of solvent, so that further dilution produces no further enthalpy change. For an ionic solid, it equals the sum of the (positive) lattice enthalpy and the (negative) enthalpy of hydration. …