Ch 1 · Periodicity of physical properties of the elements in Period 3 · AS & A Level Chemistry 9701

Topic 1 of 3 in The Periodic Table: chemical periodicity — AS & A Level Chemistry 9701 — 58 lessons.

Introduction

Periodicity is the repeating pattern of properties across a period, caused by the regular way electron configuration and nuclear charge change from one element to the next. Period 3 (sodium to argon) is the model example: every trend can be explained from structure and bonding, and Cambridge examiners expect the explanation, not just the shape of the graph.

Key ideas — Atomic radius across Period 3

Atomic radius decreases steadily from sodium to chlorine. Across the period, each successive element has one more proton in its nucleus, so nuclear charge increases by one unit at a time. The extra electrons are added to the same outer shell (n = 3), so the shielding provided by the inner shells stays essentially constant. The increasing effective nuclear charge therefore pulls the outer-shell electrons in more strongly, drawing them closer to the nucleus and decreasing atomic radius at each step.

🧑‍🏫 Teacher's explanation

Imagine a game of tug-of-war where the nucleus is the center and the electrons are the players. In Period 3, as you move from Sodium to Chlorine, the nucleus gets stronger because it adds one proton at a time. However, the number of 'shielding' layers (the inner electrons) stays exactly the same.

📖 Definitions

🔍 Worked example — Easy

Compare the atomic radius of Sodium (Z=11) and Magnesium (Z=12).

🔍 Worked example — Standard

Explain why the atomic radius decreases from Phosphorus to Sulfur.

🔍 Worked example — Exam-style

An atom has a nuclear charge of +13 and its outer electrons are in the n=3 shell. Another atom has a nuclear charge of +17 and its outer electrons are also in the n=3 shell. Which has the smaller radius?

⚠️ Common errors

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🎯 Exam question

Key ideas — Ionic radius across Period 3

The ions formed by Period 3 elements are not all the same size as their parent atoms. Na⁺, Mg²⁺ and Al³⁺ each have the electron configuration [Ne] (one fewer shell than the neutral atom), so they are considerably smaller than Na, Mg and Al themselves; radius decreases from Na⁺ to Al³⁺ because increasing nuclear charge acts on the same number of electrons. By contrast, P³⁻, S²⁻ and Cl⁻ have the electron configuration [Ar], gained by adding electrons to the existing outer shell; these anions are much larger than their parent atoms because the extra electrons increase inter-electron repulsion while nuclear charge stays the same. Overall, ionic radius falls sharply from Na⁺ to Al³⁺, jumps up at P³⁻, then falls again from P³⁻ to Cl⁻ as nuclear charge increases across the isoelectronic anions.

🧑‍🏫 Teacher's explanation

Imagine a group of people standing in a circle around a central magnet. If the magnet gets stronger (more protons) but the number of people (electrons) stays the same, the people will be pulled closer to the center. This is exactly what happens when we move from Na^+ to Al^3+. Even though they have the same number of electrons, the Al^3+ nucleus has 13 protons while Na^+ only has 11, so the pull is much stronger, shrinking the ion.

📖 Definitions

🔍 Worked example — Easy

Explain why the ionic radius of Mg^2+ is smaller than the atomic radius of Mg.

🔍 Worked example — Standard

Compare the ionic radii of Na^+ and Al^3+. Which is larger and why?

🔍 Worked example — Exam-style

Predict the trend in ionic radii for the anions P^3-, S^2-, and Cl^-.

⚠️ Common errors

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🎯 Exam question

Key ideas — Structure and bonding across the period

Na, Mg and Al are metals with giant metallic structures: positive ions in a lattice surrounded by a sea of delocalised outer-shell electrons. Si has a giant covalent (macromolecular) structure, with every atom joined to four neighbours by strong covalent bonds, similar to diamond. P (as P₄), S (as S₈), Cl (as Cl₂) and Ar (monatomic) are simple molecular or atomic structures, held together in the solid only by weak instantaneous dipole–induced dipole (van der Waals) forces between molecules.

🧑‍🏫 Teacher's explanation

Imagine a crowded ballroom. In the first group of elements (metals), the dancers (nuclei) are arranged in a strict grid, but they are surrounded by a massive, flowing crowd of people (delocalised electrons) that can move anywhere. This 'sea' is why metals conduct electricity.

📖 Definitions

🔍 Worked example — Easy

Explain why the melting point of Magnesium (Mg) is higher than that of Sodium (Na).

🔍 Worked example — Standard

Compare the structures and melting points of Silicon (Si) and Sulfur (S_8).

🔍 Worked example — Exam-style

Predict the trend in atomic radius from Na to Cl and explain it.

⚠️ Common errors

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🎯 Exam question

Key ideas — Melting point trend across Period 3

Melting point rises from Na to Al, reaches a sharp maximum at Si, then falls sharply to P, S, Cl₂ and Ar. From Na to Al, ionic charge increases (+1, +2, +3) and cationic radius decreases, so the electrostatic attraction between the metal cations and the delocalised electrons strengthens, increasing the melting point. Silicon's melting point is far higher than any of the metals because melting a giant covalent lattice requires breaking many strong covalent bonds throughout the structure. From P to Ar, melting only needs to overcome weak van der Waals forces between discrete molecules, so melting points are low; among these, melting point increases with the number of electrons per molecule (S₈ > P₄ > Cl₂ > Ar), since more electrons give stronger instantaneous dipole–induced dipole forces.

🧑‍🏫 Teacher's explanation

Imagine a construction project where you are building structures with different 'glue'. In the first part of Period 3 (the metals), you are building a massive, interconnected grid of metal ions surrounded by a 'sea' of electrons. As you move from Sodium to Aluminium, the 'glue' gets stronger because the nucleus pulls the electrons in tighter, making the structure more compact and harder to pull apart.

📖 Definitions

🔍 Worked example — Easy

Explain why the melting point of Magnesium (Mg) is higher than that of Sodium (Na).

🔍 Worked example — Standard

Compare the melting points of P_4 and S_8. Which is higher and why?

🔍 Worked example — Exam-style

A substance in Period 3 has a very high melting point and is a poor conductor of electricity in the solid state. Identify the element and explain its properties.

⚠️ Common errors

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🎯 Exam question

Key ideas — Electrical conductivity trend across Period 3

Na, Mg and Al conduct electricity well, because their delocalised electrons are free to move through the metallic lattice and carry charge; conductivity increases from Na to Al as each atom contributes more delocalised electrons per ion. Si conducts only very slightly (it is a semiconductor): its giant covalent structure has electrons largely localised in covalent bonds, with very few free to move. P, S, Cl₂ and Ar do not conduct at all in any state, since they have no delocalised electrons or mobile ions.

🧑‍🏫 Teacher's explanation

Imagine a crowded ballroom. In a metallic structure (like Na or Mg), the dancers (electrons) are not tied to any single person; they move freely across the entire floor, allowing a current to flow easily. As you move across Period 3, the 'dancers' become more tightly held by their partners (the nuclei), making it harder for them to roam.

Ch 1 · Periodicity of physical properties of the elements in Period 3 · AS & A Level Chemistry 9701

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