Valence electron

Valence electrons are the outermost electrons of an atom, and they can form chemical bonds if that outer shell isn’t full. In a single covalent bond, each atom contributes one valence electron to a shared pair. These electrons largely determine an element’s chemical behavior—its valence, how readily it bonds, and with how many other atoms. Reactivity depends heavily on the atom’s electron configuration.

For main-group elements, valence electrons are only in the highest principal energy level. For transition metals, they can also come from an inner shell. An atom with a full valence shell (like a noble gas) is chemically inert. Atoms with one or two extra valence electrons beyond a closed shell are highly reactive because those electrons are easy to remove, forming a positive ion. Atoms missing one or two electrons from a closed shell are also reactive, tending to gain or share electrons to fill the gap.

Like core electrons, valence electrons can absorb or release energy as photons. Gaining energy can excite an electron to a higher shell, or even ionize the atom by freeing the electron. Losing energy (emitting a photon) lets an electron move to an unfilled inner shell.

The highest-energy electrons determine an atom’s chemical reactivity. For phosphorus, the configuration 1s² 2s² 2p⁶ 3s² 3p³ gives five valence electrons (3s² 3p³), matching its maximum valence of 5 in PF₅. This is often written as [Ne] 3s² 3p³. For transition metals, the (n−1)d and ns energy levels are very close, so d electrons outside the noble-gas core also count as valence electrons. Manganese, with configuration [Ar] 4s² 3d⁵, has seven valence electrons, consistent with its +7 oxidation state in permanganate. However, having that many valence electrons doesn’t guarantee that oxidation state exists—fluorine never reaches +7, and the highest known oxidation state for any element is +9, despite ytterbium and nobelium having up to 16 valence electrons.

In each transition metal series, as you move right, d subshell electrons drop in energy and become less valence-like. Nickel has ten valence electrons in principle (4s² 3d⁸), but its oxidation state never exceeds +4. For zinc, the 3d subshell is always full in compounds, though it can contribute to the valence band. Similar patterns hold for the (n−2)f levels in inner transition metals. The d electron count is another way to understand transition metal chemistry.

The number of valence electrons can be read from the periodic table group. For groups 1–12, the group number equals the number of valence electrons; for groups 13–18, the units digit does. Helium is an exception: it has two valence electrons like the alkaline earth metals, but its shell is full, making it inert, so it’s placed in group 18.

The valence shell is the set of orbitals energetically available for bonding. For main-group elements, it’s the ns and np orbitals of the outermost shell. For transition metals, it includes the incomplete (n−1)d subshell; for lanthanides and actinides, the incomplete (n−2)f and (n−1)d subshells. These orbitals may come from different principal quantum numbers but are similar in energy. Generally, main-group elements (except hydrogen and helium) react to achieve an s²p⁶ configuration—the octet rule, where each bonded atom has eight valence electrons including shared ones. Transition metals tend toward a d¹⁰s²p⁶ configuration.