Periodic trends: how atomic radius, electronegativity and ionization energy change

Two ideas behind every trend

Almost every trend comes from two effects. The first is nuclear charge. Across a period, each element has one more proton, so the nucleus pulls harder on the electrons. Electrons added to the same shell do little to shield one another, so the outer electrons feel a stronger pull, called the effective nuclear charge, from left to right.

The second is the number of shells. Down a group, each element has one more occupied shell, so the outer electrons sit farther from the nucleus and are shielded by more inner electrons. They are held less tightly, even though the nucleus has more protons.

Atomic radius

Atomic radius measures the size of an atom. Across a period it generally decreases, because the stronger pull draws the electron cloud inward. Down a group it increases, because each new shell sits farther out. Positive ions are smaller than their atoms, and negative ions are larger.

The alkali metals show the group trend clearly in this site’s data: lithium 182 pm, sodium 227 pm, potassium 275 pm, rubidium 303 pm and cesium 343 pm. Radius values depend on how they are measured (covalent, metallic or van der Waals radii), so different tables give different numbers, and the trend across a period is a general pattern rather than a perfectly smooth one.

Electronegativity

Electronegativity measures how strongly an atom attracts the shared electrons in a chemical bond. On the Pauling scale, it rises across a period and falls down a group.

In period 2 it climbs from lithium (0.98) through carbon (2.55) and oxygen (3.44) to fluorine (3.98), the most electronegative element. Down group 17 it falls from fluorine to chlorine (3.16), bromine (2.96) and iodine (2.66). The lowest values are at the lower left, such as cesium at 0.79. Helium, neon and argon have no value because they form no stable compounds.

The difference in electronegativity predicts the type of bond. A large difference, as between sodium (0.93) and chlorine (3.16), gives an ionic bond, while a small difference gives a covalent bond.

Ionization energy

Ionization energy is the energy needed to remove the outermost electron from an atom in the gas phase. It rises across a period, because the electrons are held more tightly, and falls down a group, because the outer electron is farther from the nucleus.

Across period 3 it rises from 5.139 eV for sodium to 15.76 eV for argon, and down group 1 it falls from 5.392 eV for lithium to 3.894 eV for cesium. Helium has the highest first ionization energy of any element, 24.587 eV. There are small dips along the way: boron is lower than beryllium because its outer electron is in a higher-energy 2p subshell, and oxygen is lower than nitrogen because its fourth 2p electron shares an orbital, where repulsion makes it easier to remove.

Metallic character and reactivity

Metallic character describes how easily an atom loses electrons. It follows the opposite pattern to ionization energy: it decreases across a period and increases down a group. That is why the most metallic elements, such as cesium and francium, sit at the lower left, and the most nonmetallic, such as fluorine and oxygen, sit at the upper right.

Reactivity follows from these trends but depends on the group. Alkali metals become more reactive down the group because their single valence electron is easier to lose: lithium reacts steadily with water, while potassium reacts violently. Halogens become less reactive down the group because larger atoms attract an extra electron less strongly, so fluorine is the most reactive halogen.