Every Enthalpy Change Animated & Explained IN UNDER 20 MINUTES

An enthalpy change is the amount of heat energy released or taken in per mole of substance during a physical or chemical change, measured in kilojoules per mole. Standard enthalpy changes are measured under standard conditions: 1 bar (100 kPa) pressure, 298 Kelvin (25°C) for A-level chemistry, and 1 mole per decimeter cubed for solutions. Standard states refer to the most stable form of a substance under these conditions (e.g., H2 gas, O2 gas, H2O liquid).

Enthalpy changes can be exothermic (heat energy is given out, surroundings' temperature increases, negative enthalpy change) or endothermic (heat energy is taken in, surroundings' temperature decreases, positive enthalpy change). This is due to bond breaking requiring energy (endothermic) and bond formation releasing energy (exothermic). In a reaction, if less energy is taken in to break bonds than is given out to form bonds, it's exothermic. If more energy is taken in to break bonds than is given out to form bonds, it's endothermic.

Mean bond enthalpy is the average enthalpy change when one mole of gaseous covalent bonds is broken. This process is always positive (endothermic). The value for forming bonds is the same but negative (exothermic). It's an average because bond strength can vary slightly between molecules due to different environments. A larger value indicates a stronger covalent bond.

The enthalpy of reaction is the enthalpy change for a specific reaction when equation quantities of reactants react. Its value (e.g., negative for an exothermic reaction) refers to the given stoichiometry. For example, doubling the reactant quantities doubles the enthalpy change.

The enthalpy of formation is the enthalpy change when one mole of a compound is formed from its constituent elements in their standard states. The equation must always show the formation of exactly one mole of the compound. It can be endothermic or exothermic, depending on the balance between energy required for bond breaking in elements and energy released from bond formation in the product.

The enthalpy of combustion is the enthalpy change when one mole of a substance is completely burnt in oxygen. It is always negative (exothermic) because burning releases heat. This is due to the strong bonds formed in products like carbon dioxide and water, releasing more energy than is absorbed to break bonds in reactants.

The enthalpy of neutralization is the enthalpy change when one mole of water is formed from the reaction between an acid and an alkali. It is always negative (exothermic) because H+ ions and OH- ions, already present from dissociated acid and alkali, combine to form water, a bond-forming process that releases heat without significant bond breaking.

The first ionization energy is the enthalpy change when one mole of gaseous atoms each loses one electron to form one mole of gaseous ions with a 1+ charge. It is always positive (endothermic) as energy is required to remove an electron. The second ionization energy is when one mole of gaseous 1+ ions each loses one electron to form one mole of gaseous 2+ ions. It is also positive and even more so than the first, due to the increased attraction from the already positive ion.

The first electron affinity is the enthalpy change when one mole of gaseous atoms each gains an electron to form one mole of gaseous ions with a 1- charge. It is usually negative (exothermic) due to the attractive force formed. The second electron affinity is when one mole of gaseous 1- ions each gains an electron to form one mole of gaseous 2- ions. It is always positive (endothermic) because energy is required to overcome the repulsion between the incoming electron and the negative ion.

The enthalpy of fusion is the enthalpy change when one mole of a solid melts to form a liquid. It is always positive (endothermic) as energy is needed to break intermolecular bonds. The enthalpy of vaporization is when one mole of a liquid vaporizes to form a gas, also always positive (endothermic). The enthalpy of sublimation is when one mole of a solid sublimes directly to a gas, also always positive (endothermic). These are all examples of the enthalpy of transition.

The enthalpy of solution is the enthalpy change when one mole of a solute dissolves in water to give an infinitely dilute solution. It can be positive or negative, depending on the energy balance between breaking solute and water bonds (endothermic) and forming attractive forces between solute and water (exothermic). The enthalpy of hydration (or hydration enthalpy) is the enthalpy change when one mole of gaseous ions dissolves in water to give an infinitely dilute solution. It is always negative (exothermic) as it only involves the formation of attractive forces.

The enthalpy of atomization is the enthalpy change when one mole of gaseous atoms is formed from a substance in its standard state. It is always endothermic as it involves breaking bonds. It differs from mean bond dissociation enthalpy in that it specifically forms one mole of gaseous atoms from the standard state, and can apply to substances without covalent bonds (like metals). For chlorine, atomization involves producing one mole of gaseous atoms from half a mole of Cl2, whereas bond dissociation refers to breaking one mole of Cl-Cl bonds to form two moles of Cl atoms.

The lattice formation enthalpy is the enthalpy change when one mole of a solid ionic lattice is formed from its constituent ions in the gas phase. It is always negative (exothermic) due to the formation of ionic bonds. The lattice dissociation enthalpy is the opposite, the enthalpy change when one mole of a solid ionic lattice is broken down into its constituent ions in the gas phase, and is always positive (endothermic) due to the breaking of ionic bonds.