The Periodic Table

Arrangement of elements

Elements are arranged in order of increasing proton (atomic) number, one after the next. Do that and chemistry repeats in vertical columns.

Groups and periods

Read the table left to right and each element has one more proton than the last — that much is by design. What makes the table useful is that elements with the same number of outer electrons end up in the same vertical column, a group, and behave chemically like each other: sodium behaves like lithium, potassium like both. Horizontal rows are periods; elements in the same period have the same number of electron shells.

Reading position from electronic configuration

Metals, non-metals and predictions

A zig-zag line running down the right side of the table separates metals (left) from non-metals (right). Metals lose electrons to form positive ions; non-metals gain them to form negative ones. Group I metals lose one electron → 1+ ion; Group II lose two → 2+; Group VII non-metals gain one → 1−; Group VI gain two → 2−. Any element’s charge in an ionic compound comes straight from its group number.

Group I — the alkali metals

The Group I metals are so reactive they can’t be left in air. They tarnish in seconds, and in water they fizz, skate across the surface and produce an alkaline solution — hence the name.

Physical properties

Group I metals are soft (softer down the group — potassium can be cut with a knife), have low densities (lithium floats on water) and unusually low melting points for metals. A freshly cut surface is shiny silver but dulls within seconds as it oxidises.

Reaction with water

All Group I metals react with water in the same pattern, producing a metal hydroxide (which dissolves to an alkaline solution) and hydrogen gas: 2M(s) + 2H₂O(l) → 2MOH(aq) + H₂(g). The vigour changes markedly down the group. Lithium fizzes gently and moves around the surface; sodium reacts more violently, melting into a shiny ball that moves rapidly; potassium ignites, burning with a lilac flame.

ExtendedPredicting Rb and Cs

The trend continues predictably. Rubidium reacts very violently, sparking as the released hydrogen ignites; caesium can explode almost instantly. The reason is atomic structure: each element down the group has one more electron shell, so its outer electron sits further from the nucleus and is more weakly held. Losing that electron — the essential step of any Group I reaction — becomes easier with each step down.

Group VII — the halogens

The halogens are the most reactive non-metals. All are diatomic, all form 1− ions, and their physical states shift from gas to liquid to solid down the group — every trend a direct consequence of atomic size.

Physical properties

Density and melting/boiling point both increase down the group: heavier atoms in bigger molecules take more energy to pull apart. The colour also darkens down the group.

Chemical trend: reactivity down the group

The halogens react by gaining one electron to fill their outer shell. In smaller atoms (fluorine, chlorine) the outer shell is close to the nucleus and the incoming electron is pulled in strongly; in larger atoms (bromine, iodine) it is further out and the pull is weaker. So reactivity decreases down the group: fluorine is the most reactive non-metal in the entire Periodic Table; iodine is comparatively sluggish.

Astatine, at the very bottom, is predicted to be an even less reactive dark solid. It is radioactive and hard to study in bulk — the predictions come from extrapolating the observed trend.

Halogen displacement reactions

The reactivity order of the halogens shows up clearly in one experiment. Add a solution of a more reactive halogen to a salt of a less reactive one, and the more reactive halogen takes the halide’s place.

The pattern

Chlorine is more reactive than bromine, which is more reactive than iodine. Any halogen higher up the group displaces the halide of one lower down from its aqueous salt: Cl₂ + 2KBr → 2KCl + Br₂ (colourless → orange-brown); Cl₂ + 2KI → 2KCl + I₂ (colourless → dark brown); Br₂ + 2KI → 2KBr + I₂. The reverse doesn’t happen — adding iodine to potassium bromide produces no reaction, because iodine is less reactive than bromine.

ExtendedReading the redox

Take Cl₂ + 2KBr → 2KCl + Br₂. Chlorine gains electrons to become Cl⁻, so chlorine is reduced; the bromide ion loses electrons to become Br₂, so it is oxidised. The chlorine is the oxidising agent; the bromide is the reducing agent. Fluorine, being even more reactive than chlorine, would displace any other halide — but it is not used in schools because of its extreme reactivity.

Transition elements

The block of metals sitting between the two main-group columns behaves like nothing else on the table. Their compounds are coloured, their catalytic reach is enormous, and they change oxidation state with ease.

Physical properties

Coloured compounds

Catalysts & variable oxidation states

Transition elements are the go-to industrial catalysts: iron for the Haber process, vanadium(V) oxide for the Contact process, nickel for hydrogenating vegetable oils into margarine, platinum in catalytic converters. The same underlying feature — the ability to switch oxidation state easily — is what makes them useful: a catalyst that can accept and release electrons cheaply can lower an activation-energy barrier.

Noble gases

The far-right column is home to the least-reactive elements in existence. What sets them apart is a full outer electron shell — the arrangement every other atom is trying to achieve.

The full-shell rule

Helium has 2 outer electrons — a full first shell. Every other noble gas has 8 outer electrons, filling its outermost shell. This arrangement is remarkably stable: no partly-filled orbitals eager to react, no easy handle for another atom to grab. They also don’t bond with themselves — while every other elemental gas is diatomic (O₂, N₂, Cl₂), the noble gases exist as single atoms: monatomic.

Physical trend

Down the group, atoms get bigger and heavier, so boiling points rise and densities increase. Helium is the second-lightest gas known (after hydrogen); radon is dense and radioactive.

Uses

In every case the pattern is the same: an inert atmosphere is needed to stop something reacting, and a noble gas is the perfect atmosphere — it simply refuses to take part.