Periodic Table revision, GCSE & A-Level Chemistry

What is the periodic table?

The periodic table arranges all known elements in order of increasing atomic number (the number of protons in the nucleus). Elements with similar chemical properties appear in the same vertical column (group), while each horizontal row (period) corresponds to one shell of electrons being filled.

The table was developed from Dmitri Mendeleev's 1869 arrangement, which left gaps for undiscovered elements and predicted their properties. Modern tables are ordered by atomic number, not atomic mass, resolving the anomalies that puzzled early chemists.

• Ordered by increasing atomic number (proton number).
• Groups (columns): elements share the same number of outer-shell electrons.
• Periods (rows): elements share the same number of occupied electron shells.
• You are given a periodic table in GCSE Chemistry exams (AQA, Edexcel, OCR).

Groups and periods

A group is a vertical column. Elements in the same group have the same number of electrons in their outer shell, so they have similar chemical properties. Group number (1-8 or 1-18) tells you the number of outer electrons for the main-group elements.

A period is a horizontal row. Moving across a period, the number of protons increases while the number of occupied shells stays the same. The outer electrons are pulled closer to the nucleus, so atomic radius decreases across a period.

• Group 1: 1 outer electron, reactive metals.
• Group 7: 7 outer electrons, reactive non-metals.
• Group 0 (8): full outer shell, unreactive noble gases.
• Atomic radius decreases across a period (more protons, same shell, more attraction).
• Atomic radius increases down a group (more shells, more shielding).

Group 1: alkali metals

The alkali metals (lithium, sodium, potassium, rubidium, caesium) are soft, low-density metals that react vigorously with water to produce a metal hydroxide and hydrogen gas. They have one electron in their outer shell, which they lose easily to form a +1 ion.

Reactivity increases down the group: the outer electron is further from the nucleus and more shielded by inner shells, so the attraction is weaker and the electron is lost more easily. This trend is the classic GCSE example of explaining reactivity using electron structure.

• Reaction with water: metal + water → metal hydroxide + hydrogen.
• Example: 2Na + 2H₂O → 2NaOH + H₂.
• Reactivity increases down the group (outer electron easier to lose).
• All Group 1 metals are stored under oil to prevent reaction with air and moisture.

Group 7: halogens

The halogens (fluorine, chlorine, bromine, iodine) are non-metals with seven outer electrons. They gain one electron to form a -1 ion, achieving a full outer shell. Reactivity decreases down the group: the outer shell is further from the nucleus with more shielding, so gaining an electron is harder.

A more reactive halogen will displace a less reactive one from a solution of its salt. This is a displacement reaction. For example, chlorine displaces bromine from potassium bromide solution because chlorine is more reactive than bromine.

• Displacement: chlorine + potassium bromide → potassium chloride + bromine.
• Reactivity decreases down the group (harder to gain an electron).
• At GCSE you may be asked to predict displacement outcomes from halogen colour changes.
• At A-Level: trend in boiling point increases down the group (stronger London forces).

Group 0: noble gases

The noble gases (helium, neon, argon, krypton, xenon) have a full outer shell of electrons, so they are very unreactive. They exist as single atoms (monatomic) and are used where an inert atmosphere is needed, for example argon in light bulbs and helium in balloons.

At A-Level, the very low boiling points and the trend in boiling point down the group (increasing with more electrons, stronger London dispersion forces) can be tested.

• Full outer shell: do not need to gain, lose or share electrons.
• Monatomic, inert, colourless gases at room temperature.
• Boiling point increases down the group (stronger intermolecular forces).

Periodicity across Period 3 (A-Level)

Periodicity is the repeating pattern of properties across periods. Across Period 3 (Na to Ar), atomic radius decreases because nuclear charge increases while shielding stays roughly constant, pulling the outer shell closer.

First ionisation energy generally increases across the period, with two small dips: at aluminium (the 3p electron is slightly higher in energy and easier to remove than the 3p paired electron of magnesium) and at sulfur (the 3p electron being removed is paired, so there is electron-electron repulsion).

Melting point peaks at silicon (giant covalent structure), is high for Na, Mg, Al (metallic bonding), drops sharply at the non-metals (simple molecular: P₄, S₈, Cl₂, Ar), and is lowest for argon (monatomic).

• Atomic radius: decreases across Period 3.
• First ionisation energy: generally increases, with dips at Al and S.
• Melting point: peaks at silicon (giant covalent), drops for simple molecular non-metals.
• Periodicity is explained by changes in bonding and structure, not just electron count.

Transition metals (A-Level)

Transition metals are d-block elements that form at least one stable ion with a partially filled d sub-shell. Their characteristic properties include variable oxidation states, the formation of coloured ions, catalytic activity, and the formation of complex ions with ligands.

Variable oxidation states arise because the 4s and 3d electrons are similar in energy, so different numbers can be involved in bonding. Coloured ions result from electron transitions between split d orbitals when ligands bond to the central metal ion. Common catalysts include iron in the Haber process and vanadium(V) oxide in the Contact process.

• Variable oxidation states: e.g. Fe²⁺ and Fe³⁺, Cu²⁺ and Cu⁺.
• Coloured ions: [Cu(H₂O)₆]²⁺ is blue, [Fe(H₂O)₆]²⁺ is pale green.
• Catalysis: Fe in Haber process, V₂O₅ in Contact process.
• Complex ions: a central metal ion surrounded by ligands (e.g. [Cu(NH₃)₄(H₂O)₂]²⁺).

Ionisation energy trends

First ionisation energy is the energy needed to remove one electron from each atom in one mole of gaseous atoms. It generally increases across a period (more protons, same shell, stronger nuclear attraction) and decreases down a group (more shielding, outer electron further from nucleus).

Successive ionisation energies increase because each electron is removed from an increasingly positive ion. Large jumps between values provide evidence for electron shells: a big jump after removing the outer-shell electrons shows that the next electron comes from an inner shell, closer to the nucleus and harder to remove.

• Increases across a period: more nuclear charge, same shell, stronger attraction.
• Decreases down a group: more shells, more shielding, outer electron further away.
• Dips at Al and S (Period 3) explained by sub-shell structure and electron pairing.
• Successive ionisation energies: large jumps mark shell boundaries.