The Periodic Table and Groups

GCSE Chemistry · Periodic Table

The periodic table

The periodic table arranges all the elements in order of increasing atomic number (number of protons). It is organised so that elements with similar properties line up, making it a powerful tool for predicting behaviour.

Groups = the vertical columns. Elements in the same group have the same number of electrons in their outer shell, which gives them similar chemical properties. The group number = number of outer-shell electrons.
Periods = the horizontal rows. Across a period, each element has one more proton and one more electron.
Metals are on the left and centre; non-metals are on the right. A "staircase" line separates them.

How the table developed

Mendeleev arranged elements by properties and left gaps for undiscovered elements, even swapping some out of strict mass order so they fit their group. His predictions were later confirmed. Modern tables are ordered by atomic number, which fixed the anomalies.

Group 1 — the alkali metals

Elements like lithium, sodium, potassium. They have 1 electron in their outer shell.

Very reactive metals; soft; react vigorously with water to produce hydrogen and an alkaline hydroxide.
Reactivity increases down the group — the outer electron is further from the nucleus, so it's lost more easily.

Group 7 — the halogens

Elements like fluorine, chlorine, bromine, iodine. They have 7 electrons in their outer shell and form −1 ions or covalent bonds.

Exist as diatomic molecules (Cl₂, Br₂...).
Reactivity decreases down the group — the outer shell is further from the nucleus, so it's harder to gain an electron.
A more reactive halogen will displace a less reactive one from its salt solution.

Group 0 — the noble gases

Helium, neon, argon, etc. They have a full outer shell (2 or 8 electrons), so they are very unreactive (inert) and exist as single atoms.

Metals vs non-metals

Metals: lose electrons to form positive ions; conductive; shiny; malleable.
Non-metals: gain/share electrons; often form negative ions or covalent bonds; poor conductors (except graphite).

Worked example

Why does reactivity increase down Group 1 but decrease down Group 7?

Group 1 atoms lose their outer electron — further down, it's further from the nucleus and lost more easily → more reactive.
Group 7 atoms gain an electron — further down, the outer shell is further away, so it's harder to attract one → less reactive. ✓

Common mistakes

Saying reactivity increases down every group — it increases down Group 1 but decreases down Group 7.
Forgetting group number = outer-shell electrons.
Thinking noble gases form ions — they're inert (full outer shell).

Exam tips

Learn the trends for Groups 1, 7 and 0 and be able to explain them using electron distance from the nucleus.
Remember Mendeleev left gaps and predicted properties.
Halogen displacement: more reactive displaces less reactive.

Key facts to remember

Ordered by atomic number; groups = columns (same outer electrons, similar properties); periods = rows.
Group 1: reactive metals, reactivity increases down; Group 7: halogens, reactivity decreases down; Group 0: inert (full outer shell).
Metals lose electrons (left/centre); non-metals gain/share (right).