How Electron Configurations Build Periods and Groups

The arrangement of electrons explains the arrangement of elements

Atomic number tells us how many electrons a neutral atom contains. Chemistry depends strongly on how those electrons are arranged, especially the outer electrons that interact with other atoms. As electrons fill a recurring sequence of shells and subshells, similar outer arrangements appear again. That is the microscopic reason for periodicity.

An orbital is a permitted electron state, not a little circular path. Each orbital can contain at most two electrons. An s subshell has one orbital, p has three, d has five, and f has seven. Therefore their capacities are 2, 6, 10, and 14 electrons respectively. These numbers explain the widths of the periodic-table blocks.

The allowed values of $l$ in a shell run from 0 to $n-1$. However, a period does not simply fill every subshell with the same $n$. Subshell energies overlap: 4s begins filling before 3d, and 6s before 4f. Use the filling sequence as well as the shell labels.

Building the periods

The first period starts with the lowest-energy 1s orbital. Hydrogen contributes one electron and helium completes the 1s pair. The next element must begin a new shell, giving lithium the configuration $1s^{2}2s^1$. That new start begins the second period.

The third period has eight elements even though a 3d subshell is allowed by the quantum numbers. The 3d subshell is not filled at that stage in the sequence. After argon, potassium and calcium begin 4s filling; the 3d series then appears before the period finishes by filling 4p. Scandium begins the 3d series with a configuration commonly written $[Ar]3d^{1}4s^2$, and zinc has $[Ar]3d^{10}4s^2$.

The fifth period repeats the s–d–p pattern using 5s, 4d, and 5p. The sixth and seventh add f-subshell filling. In the layout used here, the fourteen detached entries Ce–Lu and Th–Lr fit into those longer rows. Actual d- and f-element ground states can depart from a simple filling prediction because subshell energies are close; the sequence explains the overall architecture.

Why groups behave like families

Once a period reaches a closed outer arrangement, the next period starts a similar outer pattern at a greater distance from the nucleus. Lithium, sodium, and potassium all have one electron in an outer s subshell. Their full configurations differ, but the shared outer pattern helps explain their common tendency to form singly positive ions.

The noble-gas symbol in square brackets represents the filled inner core. Thus $[Ne]3s^1$ means the neon electron arrangement plus one electron in 3s. Members of Group 1 share the general outer configuration $n{s}^1$. Their chemistry is related, but their sizes and ease of electron removal change down the group, so related behaviour does not mean identical reactivity.

For main-group elements, the largest occupied principal quantum number normally gives the period. An outer configuration $n{s}^1$ identifies Group 1 and $n{s}^2$ identifies Group 2, apart from helium’s special position. Configurations $n{s}^{2}n{p}^1$ through $n{s}^{2}n{p}^6$ identify Groups 13 through 18. For these p-block elements, add 10 to the number of outer s and p electrons to obtain the group number.