Ionization enthalpy

Ionisation enthalpy is the energy needed to remove the most loosely held electron from an isolated gaseous atom in its ground state: X(g) → X⁺(g) + e⁻. Since energy must always be supplied to pull an electron away, it is always positive.

Electrons can be removed one after another, and each successive ionisation enthalpy is larger than the one before, because the remaining electrons are held more tightly by a nucleus that now outweighs them. It is like students leaving a class with a strict teacher: the first slips out easily, the second has a harder time, the third harder still. A very big jump shows that the next electron must come from a noble-gas core. Sodium's first value is 496 kJ mol⁻¹, but its second is enormous, because it would break into the stable neon core.

The trends:

• Across a period, ionisation enthalpy increases: the nuclear charge grows while the electrons stay in the same shell.
• Down a group, it decreases: the atoms are bigger, and the inner electrons shield the outer one from the nucleus.

Learn the general trend first, then the exceptions:

• Boron is lower than beryllium: boron's outer electron is in 2p, which is shielded by the 2s electrons and held less strongly.
• Oxygen is lower than nitrogen: nitrogen's half-filled 2p³ subshell is extra stable; in oxygen, one 2p orbital holds a pair of electrons, which repel each other, so one is easier to remove.
• In group 13, gallium is about the same as aluminium, and thallium is higher than indium, because their inner d and f electrons shield poorly.

Isotopes of an element have the same electronic configuration and nuclear charge, so they have the same ionisation enthalpy.