Enthalpy of Atomisation (ΔaH°) is defined as the enthalpy change when one mole of a substance is completely converted into gaseous atoms in its standard state. Since breaking bonds always requires energy input, the enthalpy of atomisation is always positive (endothermic).
The enthalpy of atomisation is always a positive value because energy is always required to overcome the attractive forces holding atoms together. No substance spontaneously breaks apart into gaseous atoms.
The process of atomisation involves breaking all the bonds in a substance until you are left with individual, isolated gaseous atoms:
For a diatomic molecule: H₂(g) → 2H(g), ΔaH° = +436 kJ/mol
This means 436 kJ of energy must be supplied to break all the H–H bonds in 1 mole of H₂ gas to produce 2 moles of isolated H atoms.
For metals: Na(s) → Na(g), ΔaH° = +108 kJ/mol For metals, atomisation = sublimation (converting solid metal directly to gaseous atoms).
The enthalpy of atomisation directly reflects the strength of bonds holding atoms together:
In transition metals, it also reflects the strength of metallic bonding, which varies across and down the periodic table.
For diatomic molecules (like H₂, Cl₂), they are the same. But for polyatomic molecules (like CH₄), the enthalpy of atomisation involves breaking ALL bonds, while bond dissociation enthalpy refers to breaking ONE specific bond at a time.
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