Ionization Energy and Electron Affinity

Which atoms give up electrons easily?

Drop a small piece of sodium into water (a teacher demonstration only, behind a safety shield) and it fizzes, melts into a ball and races around the surface. Potassium does the same thing even more violently. Neon, in a glowing sign, does not react with anything at all.

These differences come down to one question: how much energy does it take to pull an electron off an atom? That quantity, ionization energy, can be measured very precisely, and its pattern across the periodic table is one of the clearest pieces of evidence that electrons are arranged in shells.

Ionization energy

The first ionization energy (IE₁) is the energy needed to remove the most loosely held electron from one mole of atoms in the gas state:

X(g) + energy → X⁺(g) + e⁻

It is measured in kilojoules per mole (kJ/mol). (A mole is a counting unit for very large numbers of particles; you will use it a lot in the Quantities in Chemical Reactions unit.) A high ionization energy means the electron is held tightly. A low ionization energy means it comes off easily.

Trends:

• Across a period, IE₁ generally increases. Zeff increases and the atoms get smaller, so valence electrons are held more tightly.
• Down a group, IE₁ decreases. The valence electrons are in shells farther from the nucleus, with more shielding, so they are easier to remove.

The noble gases have the highest values in each period (helium has the highest of all, 2372 kJ/mol) and the alkali metals the lowest. That is why metals on the left form positive ions so easily and the noble gases hardly react.

Small dips: B after Be, O after N

Look closely at the graph. Boron's IE₁ (801 kJ/mol) is a little lower than beryllium's (899), and oxygen's (1314) is a little lower than nitrogen's (1402). The same dips appear in period 3 (Al after Mg, S after P). These small exceptions are real data, not errors. They are explained by the way electrons are arranged in sub-shells (orbitals) within each shell, which you will study in Grade 12 Chemistry (SCH4U). For now, notice that the general trend across the period still holds.

Successive ionization energies

After one electron is removed, a second and third can be removed too, but each needs more energy, because the remaining electrons are held by the same protons with less electron repulsion.

The really useful clue is a huge jump. For magnesium (2, 8, 2), the 1st and 2nd ionization energies are fairly low (738 and 1451 kJ/mol), but the 3rd is 7733 kJ/mol, more than five times the 2nd. The third electron must come from the core (shell 2), much closer to the nucleus and not shielded by any other shell. The jump tells you magnesium has 2 valence electrons, which is exactly why magnesium forms Mg²⁺ and never Mg³⁺.

Electron affinity

The electron affinity of an element is the energy change when one mole of gaseous atoms gains one mole of electrons:

X(g) + e⁻ → X⁻(g)

For most elements, energy is released when an electron is added, because the electron is attracted by the nucleus. The more energy released, the more strongly the atom attracts an extra electron. (Some data tables write released energy as a negative number, such as −349 kJ/mol for chlorine. In this lesson we give the amount of energy released as a positive number and say "released".)

Trends (general, with more exceptions than ionization energy):

• Across a period, more energy is generally released, reaching a maximum at the halogens (group 17). A halogen needs only one electron to complete its outer shell, and its large Zeff pulls that electron in strongly.
• Down a group, less energy is generally released, because the added electron goes into a shell farther from the nucleus. (Fluorine is an exception: it releases a little less than chlorine, because its very small outer shell is crowded and the electrons repel the newcomer.)
• The noble gases do not form stable negative ions: an added electron would have to start a new shell, far from the nucleus.

Putting it together: metals and non-metals

• Metals (left side) have low ionization energies and small electron affinities. They lose electrons easily and form cations.
• Non-metals (right side, except noble gases) have high ionization energies and large electron affinities. They gain electrons and form anions.
• Reactivity: for metals, the lower the ionization energy, the more reactive the metal. That is why reactivity increases down group 1 (Li < Na < K). For halogens, reactivity decreases down the group (F > Cl > Br > I), because larger atoms attract electrons less strongly.