The periodic table for A-Level Chemistry AQA

The periodic table is the chemist's most powerful pattern-finder. At AQA A-Level Chemistry, it is the foundation for unit 3.2.1, where you learn how elements are arranged by atomic number into groups and periods that repeat predictable chemical and physical behaviour. Once you can read the table, you can predict how almost any element will behave.

This guide covers the AQA specification for periodicity: How the table is built, the four big trends across Period 3 (atomic radius, melting point, first ionisation energy, electronegativity), and the underlying reasons examiners want you to explain.

The modern periodic table arranges 118 confirmed elements by atomic number (the number of protons in the nucleus). Elements are placed in horizontal rows called periods and vertical columns called groups. Elements in the same group share similar chemical properties because they have the same number of electrons in their outer shell.

The table is divided into four blocks based on which sub-shell the outermost electron occupies. The s-block holds Groups 1 and 2, the p-block holds Groups 3 to 0 (excluding the transition metals), the d-block holds the transition metals (Groups 3 to 12), and the f-block holds the lanthanides and actinides.

Knowing the block tells you the highest-energy sub-shell the element is filling. AQA examiners often phrase electron configuration questions as "Which block is element X in?" and you need to answer using sub-shell notation.

Atomic radius decreases across a period and increases down a group. Across Period 3 (Na to Ar), the radius shrinks because each successive element has one more proton in the nucleus but no extra shell of electrons. The increased nuclear charge pulls the existing electrons in tighter, even though shielding stays roughly the same.

Down a group, the radius grows because each new period adds a new electron shell. The outer electrons sit further from the nucleus and experience more shielding from inner shells, so the atom is bigger.

Melting point across Period 3 is a great exam question because it tests structure as well as bonding. Metals (Na, Mg, Al) have giant metallic structures with delocalised electrons. Silicon (Si) is a giant covalent (macromolecular) network with strong covalent bonds throughout. Phosphorus, sulfur, chlorine and argon are simple molecular substances held together by weak van der Waals forces.

The melting point rises from Na to Al (more delocalised electrons per atom, stronger metallic bonding), peaks at Si (giant covalent), then drops sharply at P (simple molecular). Among the simple molecular elements, S₈ is the largest molecule and so has the highest melting point of the four.

First ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms. The trend across Period 3 is a general increase from Na to Ar, because nuclear charge increases while shielding stays roughly the same and atomic radius decreases. All three factors make the outer electron harder to remove.

There are two dips you must explain. From Mg to Al, the outer electron moves from 3s to 3p (3p is slightly higher energy and easier to remove). From P to S, paired electrons in 3p experience repulsion, so the first one removed from S is easier than expected. These two anomalies are AQA exam favourites.

Electronegativity is the power of an atom to attract a bonding pair of electrons in a covalent bond. It is measured on the Pauling scale. Across a period, electronegativity increases (more protons, smaller radius, stronger pull on bonding electrons). Down a group, it decreases (more shielding, larger radius, weaker pull).

Fluorine has the highest electronegativity (4.0). The most electronegative elements (F, O, N, Cl) are crucial for understanding hydrogen bonding, polarity and intermolecular forces later in the course. Noble gases are usually excluded because they do not normally form bonds.

Trends down a group are the reverse of those across a period for the same reason. Each step down adds a new shell, so atomic radius grows, shielding increases, and the nucleus has a weaker grip on the outer electrons. The result: Ionisation energy falls, electronegativity falls, and metallic character rises.

This is why Group 1 elements get more reactive down the group (easier to lose the outer electron) but Group 7 elements get less reactive down the group (harder to gain an electron). The same three-factor explanation works for both.

Examiner reports list a familiar pattern of errors. Most are about completeness: Stating the trend without explaining why, or quoting only one of the three factors. Always state the trend, then explain it using nuclear charge, shielding and distance.

Question: Explain why the first ionisation energy of magnesium is higher than that of aluminium.

Step 1: State the rule being broken. The general trend is that ionisation energy increases across a period.

Step 2: Identify the electron being removed. Magnesium's outer electron is in 3s; aluminium's is in 3p.

Step 3: Explain energy levels. The 3p sub-shell is at a higher energy than 3s, so the 3p electron is easier to remove.

Step 4: Add shielding. The 3p electron is slightly shielded by the 3s electrons, further reducing the energy needed.

Conclusion: Despite aluminium having a higher nuclear charge, the combination of sub-shell energy and slight extra shielding makes its first ionisation energy lower than magnesium's.