The Periodic Table

Elements are laid out in order of atomic number, but the genius of the table is the wrapping: each row (a period) fills one more electron shell, and each column (a group) collects elements with the same number of outer-shell electrons. Because chemistry is decided by outer electrons, a group is a family that behaves alike.

Metals sit on the left and centre, non-metals on the right, with a staircase of semi-metals between. The group number (for the main groups) tells you the valence electrons directly — group 1 has one, group 7 has seven.

Switch between the properties and watch the pattern: metals to the left, reactivity and size building toward the bottom-left for metals and the top-right for non-metals.

Lithium, sodium, potassium: one outer electron they lose easily, so they are very reactive metals, stored under oil and forming +1 ions. React with water to give a hydroxide and hydrogen: 2Na + 2H₂O → 2NaOH + H₂.

The key trend: reactivity increases down the group. Lower down, the outer electron is further from the nucleus and more shielded, so it is lost more easily — potassium reacts far more violently than lithium.

Fluorine, chlorine, bromine, iodine: seven outer electrons, needing just one more, so they are reactive non-metals forming −1 ions. They exist as diatomic molecules (Cl₂, Br₂).

Their trend runs the opposite way: reactivity decreases down the group, because gaining an electron gets harder as the outer shell sits further out. This gives displacement: a more reactive halogen pushes a less reactive one out of its salt — chlorine displaces bromine from potassium bromide.

Group 0 — the noble gases have a full outer shell, so they are unreactive (inert) and exist as single atoms. That stability is why they are used where you need something that will not react: argon in light bulbs, helium in balloons.

The transition metals in the centre are the "typical metals" — hard, dense, high melting points — but with extras: they form ions of more than one charge (Fe²⁺ and Fe³⁺), their compounds are often coloured, and many act as catalysts (iron in the Haber process).

The periodic table works because chemistry is decided by outer-shell electrons. Each new period starts a new shell, and across the period that shell fills up — so the same outer arrangements recur, and with them the same kinds of chemistry. The group number equals the number of outer electrons, and the period number equals the number of occupied shells.

That single idea explains everything else on the table: why Group 1 metals all react the same way, why Group 0 is inert, and why reactivity trends run the way they do.

Move across Period 3 (Na → Ar) and clear trends appear. Atomic radius decreases: the nuclear charge rises while electrons enter the same shell, pulling everything in tighter. First ionisation energy generally rises for the same reason, with small dips where a new sub-shell or a paired orbital begins.

Melting point rises to a peak at silicon then collapses: Na, Mg and Al are metallic (rising with more delocalised electrons), silicon is a giant covalent structure (very high), while P₄, S₈, Cl₂ and Ar are simple molecular or atomic with only weak forces (very low).