The Periodic Table Explained — GCSE Chemistry Group by Group

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The periodic table is one of the most powerful tools in chemistry — it arranges all known elements in a way that reveals patterns in their properties and predicts how they will behave. At GCSE, you need to know the trends within specific groups and be able to explain them in terms of atomic structure. This guide covers the groups tested most heavily in exams.

How the Periodic Table Is Organised

Elements are arranged in order of increasing atomic number (number of protons). Horizontal rows are called periods. Vertical columns are called groups. Elements in the same group have the same number of electrons in their outer shell — this is why they have similar chemical properties. Elements in the same period have the same number of electron shells.

The group number tells you the number of outer electrons. Group 1 elements have 1 outer electron. Group 7 elements have 7. Group 0 (noble gases) have a full outer shell (8 electrons, or 2 for helium).

Group 1 — The Alkali Metals

Group 1 contains lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr). All have one electron in their outer shell, which they lose easily to form 1+ ions.

Reactions with Water

All Group 1 metals react vigorously with cold water to produce a metal hydroxide and hydrogen gas. The reaction gets more vigorous down the group.

• Lithium: fizzes steadily, moves around the surface
• Sodium: fizzes vigorously, melts into a ball and moves rapidly
• Potassium: bursts into lilac/violet flame, very violent reaction

The general equation: 2M + 2H₂O → 2MOH + H₂ (where M is the Group 1 metal)

Why Reactivity Increases Down Group 1

As you go down Group 1, each element has more electron shells. The outer electron is further from the nucleus and shielded from the positive nuclear charge by more inner electron shells. This means the outer electron is held less tightly and is lost more easily. The more easily the outer electron is lost, the more reactive the element — which is why caesium is far more reactive than lithium.

Group 7 — The Halogens

Group 7 contains fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At). All have seven electrons in their outer shell and need one more to achieve a full outer shell — they gain electrons easily, forming 1− ions.

Physical Properties Down Group 7

• Fluorine: pale yellow gas at room temperature
• Chlorine: yellow-green gas at room temperature
• Bromine: orange-brown liquid at room temperature
• Iodine: dark grey solid at room temperature, purple vapour

Melting and boiling points increase down the group — molecules get larger, with more electrons, so intermolecular forces between molecules become stronger and more energy is needed to separate them.

Why Reactivity Decreases Down Group 7

Going down Group 7, each element has more electron shells. The outer shell (where the electron needs to be gained) is further from the nucleus and more shielded. Gaining an electron becomes progressively harder — the incoming electron is attracted less strongly. So fluorine is the most reactive halogen, iodine much less so.

Displacement Reactions

A more reactive halogen can displace a less reactive halogen from its salt solution. Chlorine displaces bromine and iodine. Bromine displaces iodine but not chlorine. Iodine cannot displace either.

Chlorine water added to potassium bromide solution: the solution turns orange-brown as bromine is displaced. Chlorine water added to potassium iodide solution: turns brown as iodine is displaced. These colour changes are tested directly in exam questions.

Group 0 — The Noble Gases

Group 0 contains helium, neon, argon, krypton, xenon and radon. All have full outer electron shells — they are extremely unreactive and exist as monatomic gases. Helium has 2 outer electrons (full for the first shell). All others have 8 outer electrons.

Boiling points increase down Group 0 for the same reason as Group 7 — larger atoms with more electrons have stronger intermolecular forces. Uses reflect their inertness: helium in balloons (non-flammable), argon in light bulbs (prevents filament oxidation), neon in signs (glows when electricity passes through).

The Transition Metals

The transition metals occupy the central block of the periodic table (Period 4, Groups 3–12). They include iron, copper, nickel, chromium, manganese, zinc and titanium. They have distinct properties compared to Group 1 metals:

• High melting points and densities: Much harder and denser than Group 1 metals.
• Variable oxidation states: They can form ions with different charges — iron forms Fe²⁺ and Fe³⁺, copper forms Cu⁺ and Cu²⁺.
• Coloured compounds: Their compounds are often distinctively coloured — copper compounds are blue, iron(II) compounds pale green, iron(III) compounds orange-brown, manganese compounds purple (in permanganate).
• Catalytic activity: Many transition metals and their compounds are important catalysts — iron in the Haber process, vanadium(V) oxide in the Contact process, platinum and rhodium in catalytic converters.

The AQA periodic table specification is at the AQA GCSE Chemistry specification page.