More complex patterns of the Periodic Table — Eduqas A-Level Chemistry

The periodic table shows trends in atomic radius, ionisation energy, electronegativity, and melting points across periods and down groups. These patterns arise from changes in nuclear charge, electron shielding, and sub-shell filling, enabling prediction of element properties and reactivity.

The periodic table is more than just a list of elements; it is a powerful tool for predicting chemical behaviour. Understanding the trends in atomic radius, ionisation energy, electronegativity, and melting points allows chemists to rationalise reactivity, bonding, and physical properties. These trends are governed by three key factors: nuclear charge, electron shielding, and the principal quantum shell of the outer electrons.

In WJEC A-Level Chemistry, you are expected to explain these trends across Periods 2 and 3, and down Groups 1, 2, and 7. You must also account for anomalies, such as the drop in ionisation energy between Groups 2 and 3, and between Groups 5 and 6. Mastery of these patterns is essential for understanding redox reactions, acid-base behaviour, and the properties of elements and their compounds.

This topic also links to later work on transition metals, periodicity in oxides and chlorides, and the interpretation of data from the periodic table. A solid grasp of these concepts will help you tackle exam questions that require you to predict properties or explain trends using electronic configurations.

• →Atomic radius decreases across a period due to increasing nuclear charge; increases down a group due to additional electron shells.
• →First ionisation energy generally increases across a period, with exceptions at Groups 2→3 and 5→6 due to sub-shell stability and electron pairing.
• →Electronegativity increases across a period and decreases down a group; fluorine is the most electronegative element.
• →Melting points across Period 3 show a peak at silicon due to its giant covalent structure, with metallic elements having moderate melting points and simple molecular elements having low melting points.
• →Electron shielding remains roughly constant across a period but increases down a group, affecting ionisation energy and atomic radius.

• Amphoteric behaviour of Al and Pb
• Inert pair effect in Groups 3, 4, and 5
• Donor-acceptor compounds (e.g., NH3.BF3)
• Structure and bonding in boron nitride
• Relative stability of oxidation states in Group 4
• Acid-base properties of CO2 and PbO
• Trends in bonding of Group 4 chlorides
• Disproportionation reactions of chlorine
• Origin of colour in transition metal complexes via d-orbital splitting
• Ligand exchange reactions and coordination number changes
• Heterogeneous and homogeneous catalysis mechanisms

• 💡Ensure you can draw 3D representations of octahedral and tetrahedral complexes
• 💡Practice writing ion/electron half-equations for disproportionation reactions
• 💡Be prepared to explain the splitting of d-orbitals in terms of ligand interaction
• 💡Memorize the specific colours of common transition metal ions in aqueous solution
• 💡Always quote specific data (e.g., atomic radii in pm, ionisation energies in kJ mol⁻¹) to support your explanations. This shows you have learned the values.
• 💡When explaining trends, mention all three factors: nuclear charge, shielding, and distance from nucleus. A common mistake is to only mention nuclear charge.
• 💡For 6-mark questions, structure your answer by element or by property. Use correct terminology such as 'delocalised electrons', 'van der Waals forces', and 'giant covalent lattice'.

• Confusing the inert pair effect with general group trends
• Misidentifying the coordination number in ligand exchange reactions
• Incorrectly explaining the origin of colour in transition metal complexes
• Failing to distinguish between heterogeneous and homogeneous catalysis mechanisms
• Misconception: Atomic radius increases across a period because more electrons are added. Correction: Although more electrons are added, they are in the same shell, and the increased nuclear charge pulls them closer, decreasing radius.
• Misconception: Ionisation energy always increases across a period. Correction: There are drops between Groups 2 and 3, and between Groups 5 and 6, due to changes in sub-shell and electron pairing.
• Misconception: Melting point increases steadily across a period. Correction: The trend is not monotonic; it peaks at silicon (giant covalent) and then drops sharply for non-metals.

1. 1Day 1-2: Review electronic configurations for elements in Periods 2 and 3, and Groups 1, 2, and 7. Understand how to write them and identify outer electrons.
2. 2Day 3-4: Learn the trends in atomic radius and ionisation energy. Draw graphs and annotate with explanations for each anomaly.
3. 3Day 5-6: Study electronegativity and melting point trends. Relate melting points to bonding and structure.
4. 4Day 7-8: Practice past paper questions, especially those asking to 'explain' trends. Use the mark schemes to check your use of key terms.
5. 5Day 9-10: Consolidate with active recall: cover the trends and try to reproduce them from memory. Teach a friend or record yourself explaining.

• 📋Multiple choice: Identifying the correct trend or anomaly (e.g., which element has the highest first ionisation energy?).
• 📋Short answer: Explain why a specific trend occurs (e.g., why does atomic radius decrease across a period?).
• 📋Data analysis: Given a table of values, describe and explain the trend (e.g., melting points of Period 3 elements).
• 📋6-mark extended response: Compare and contrast properties across a period or down a group, often with a focus on bonding and structure.