The Periodic Table

Core Concepts

Periods, groups, and what they encode

The table has 7 periods (horizontal rows) and 18 groups (vertical columns). The period number tells you the highest principal energy level (n) occupied in the ground-state configuration: all period-3 elements (Na through Ar) have outer electrons in the n = 3 shell.

For main-group elements (groups 1, 2, and 13–18), the group number gives the valence electron count — the electrons available for bonding. Groups 1 and 2 have 1 and 2 valence electrons. For groups 13–18:

valence electrons = group number - 10

Sulfur, in group 16, has 16 - 10 = 6 valence electrons. This single relationship explains why elements in the same group form similar compounds: they carry the same number of bonding electrons.

Metals, nonmetals, and metalloids

About three-quarters of the elements are metals — shiny, malleable, ductile solids that conduct heat and electricity. They occupy the left and center of the table. Nonmetals — often gases or brittle solids that conduct poorly — sit in the upper right, plus hydrogen. Between them lies a diagonal staircase of metalloids (B, Si, Ge, As, Sb, Te, Po, and sometimes At) with intermediate properties. Silicon's half-metallic character is exactly what makes it useful in computer chips.

Families with names

• Group 1 — alkali metals (Li, Na, K...): soft, highly reactive metals forming +1 ions.
• Group 2 — alkaline earth metals (Be, Mg, Ca...): reactive metals forming +2 ions.
• Group 15 — pnictogens (N, P, As...): the nitrogen family.
• Group 16 — chalcogens (O, S, Se...): the oxygen family.
• Group 17 — halogens (F, Cl, Br, I, At): reactive nonmetals forming −1 ions; "halogen" means salt-former.
• Group 18 — noble gases (He, Ne, Ar...): extremely unreactive, with filled valence shells.

Groups 3–12 are the transition metals, which form variable-charge ions (Fe²⁺ and Fe³⁺); the two rows pulled below the table are the lanthanides and actinides.

Blocks follow electron configurations

Each region matches the subshell being filled: s-block (groups 1–2), p-block (13–18), d-block (3–12), f-block (lanthanides and actinides). This makes shorthand configurations easy: phosphorus, 1s² 2s² 2p⁶ 3s² 3p³, is [Ne]3s²3p³. The block gives the last subshell filled, the period gives n, and the group gives the number of outer electrons.

The table as a trend map

The geometry encodes the trends from the previous topic: atomic radius increases down a group and decreases across a period; ionization energy and electronegativity generally do the reverse. You do not need to memorize 118 values — just read the arrows.

Worked Example: Finding an Element from Its Configuration

Problem. An element has the configuration 1s² 2s² 2p⁶ 3s² 3p³. Identify its period, group, block, and name.

Solution. The highest occupied level is n = 3, so it is in period 3. The outer electrons fill 3p, so it is in the p-block. Counting n = 3 electrons: 2 + 3 = 5 valence electrons. Apply the group formula:

group number = valence electrons + 10 = 5 + 10 = 15

Group 15 is the pnictogens; the period-3 member is phosphorus (P). A p-block nonmetal, it forms compounds such as phosphate, PO₄^3-, matching family behavior.

Worked Example: Predicting Ion Charge from Position

Problem. Predict the most common ion charge for magnesium (group 2) and chlorine (group 17), and write each ion's configuration.

Solution. Magnesium has two valence electrons, [Ne]3s². Losing both gives the stable neon configuration:

Mg → Mg²⁺ + 2e^-   Mg²⁺: [Ne]

Chlorine has seven valence electrons, [Ne]3s² 3p⁵. Gaining one completes its octet:

Cl + e^- → Cl^-   Cl^-: [Ne]3s² 3p⁶ = [Ar]

So Mg forms Mg²⁺ and Cl forms Cl^-. The resulting compound, MgCl₂, is the payoff of reading the table: group 2 → +2, group 17 → −1, balanced as one Mg²⁺ for two Cl⁻.

Common Confusions