Periodic Trends in Ionization Enthalpy

The ionization enthalpy ($IE$, also written $\Delta_i H$) of an element is the minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state, forming a gaseous cation:

This is more precisely called the first ionization enthalpy ($IE_1$); the energy needed to remove a second electron from the resulting $M^{+}(g)$ ion is the second ionization enthalpy ($IE_2$), and so on for successive removals. Ionization enthalpy is always a positive (energy-absorbing) quantity, since pulling a negatively charged electron away from a positively charged nucleus always requires energy input.

Trend across a period. $IE_1$ generally increases on moving left to right across a period, for the same reason atomic radius shrinks: effective nuclear charge rises while the outer electron stays in the same shell, so the nucleus holds that electron more tightly and more energy is needed to remove it. Across Period 3, $Na < Mg < Al < Si < P < S < Cl < Ar$ in $IE_1$, in general.

Trend down a group. $IE_1$ generally decreases on descending a group. The outermost electron is in a progressively higher, larger shell, farther from the nucleus and increasingly shielded from the nuclear charge by the filled inner shells, so it is held less tightly and removed more easily. Down Group 1, $Li > Na > K > Rb > Cs$ in $IE_1$.

Two classic exceptions students must know, both caused by extra-stable electron configurations resisting disruption:

• $IE_1(Be) > IE_1(B)$ — even though boron has one more proton than beryllium. Beryllium's configuration $1s^2\,2s^2$ has a completely filled $2s$ subshell, which is an unusually stable arrangement; removing an electron from it is harder than expected. Boron's extra electron instead starts a new $2p$ subshell ($1s^2\,2s^2\,2p^1$); a $2p$ electron is, on average, farther from the nucleus and less tightly bound than a $2s$ electron of the same shell, so it is removed more easily despite boron's higher nuclear charge.
• $IE_1(N) > IE_1(O)$ — even though oxygen has one more proton than nitrogen. Nitrogen's configuration $1s^2\,2s^2\,2p^3$ has a half-filled $2p$ subshell (one electron in each of the three $2p$ orbitals), which is also an unusually stable, extra-symmetric arrangement. Oxygen's extra electron ($1s^2\,2s^2\,2p^4$) must pair up with one of these three, and the resulting electron-electron repulsion within that doubly occupied orbital makes it easier to remove one electron from oxygen than expected, dropping its $IE_1$ below nitrogen's. …