Periodic table with valence electrons

What are valence electrons?

Electrons fill shells around the nucleus, and the outermost occupied shell is the valence shell. Its electrons are the farthest from the nucleus and the least tightly held, so they are the ones that are lost, gained or shared when atoms bond. Elements with the same number of valence electrons behave in similar ways, which is why they line up in the same group.

Many atoms react in ways that give them eight electrons in their outer shell, a pattern called the octet rule (hydrogen and helium are stable with two). Sodium, with 1 valence electron, readily loses it, chlorine, with 7, readily gains one more, and the noble gases, which already have a full shell, rarely react.

Finding valence electrons from the group number

Group 1 elements have 1 valence electron and group 2 elements have 2. For groups 13 to 18, subtract 10 from the group number: boron and aluminum in group 13 have 3, carbon in group 14 has 4, nitrogen 5, oxygen 6, the halogens 7 and the noble gases 8. Helium is the exception in group 18, with a full first shell of just 2.

In electron-configuration terms, a main-group element's valence electrons are those in its outer s and p subshells. Sulfur, [Ne] 3s² 3p⁴, has 2 + 4 = 6. On the table, each cell shows the number, and the legend under the table shows the color scale.

Transition metals, lanthanides and actinides

For these metals the table counts the electrons in the outermost shell, which are normally just the outer s electrons, so most show 2. Iron, [Ar] 3d⁶ 4s², is shown with 2, and chromium and copper, which each have a single 4s electron, are shown with 1. Palladium is the exception: its configuration, [Kr] 4d¹⁰, has no electron in the 5s subshell, so it is shown with its 10 outer d electrons.

Keep in mind that d electrons, and f electrons in the lanthanides and actinides, can also take part in bonding. That is why iron forms both Fe²⁺ and Fe³⁺ and manganese can reach +7. For these metals the number shown is a starting point, not a full count of the electrons they can use.

Worked example: oxygen

Oxygen has atomic number 8 and the configuration [He] 2s² 2p⁴. Its outer shell, the second, holds 2 + 4 = 6 electrons, which matches its place in group 16. To complete an octet it needs two more. It can gain them to form the oxide ion, O²⁻, or share them in two covalent bonds, as it does in water (H₂O). In the Lewis structure of water, oxygen has two bonding pairs and two lone pairs, eight electrons in total.