UNIT 2 CHEMISTRY · PEARSON AAQ · C1.1
Physical Properties Across Period 3
FREE CHEMISTRY REVISION · TOPIC 18
Explain the trends by following electrons, bonds and structures.
Connect radius, melting-point and conductivity patterns from sodium to argon to their electronic structures and the particles present.
Pearson BTEC Level 3 AAQ Applied Science
Unit 2: Principles and Applications of Chemistry · C1.1
What you will learn
Explain changes in atomic and ionic radius, melting point and electrical conductivity across Na, Mg, Al, Si, P, S, Cl and Ar. Connect the electronic structures from Section A and bonding models from Section B to the measured physical properties, rather than memorising an unexplained graph.
1. Electronic structure across the period
All Period 3 atoms have their highest occupied principal shell at n = 3. From Na to Ar the proton number rises from 11 to 18 and electrons are added to 3s and then 3p. The inner [Ne] core remains the same: Na [Ne]3s¹, Mg [Ne]3s², Al [Ne]3s²3p¹, Si [Ne]3s²3p², P [Ne]3s²3p³, S [Ne]3s²3p⁴, Cl [Ne]3s²3p⁵, Ar [Ne]3s²3p⁶.
This shared shell pattern helps explain size, but the bonding and structure change substantially. Na–Al are metallic, silicon has a giant covalent structure, white phosphorus and sulfur have discrete P₄ and S₈ molecules, chlorine is Cl₂ and argon consists of individual atoms. State which allotrope or physical form you are comparing.
2. Atomic radius: attraction, shielding and definitions
For a consistent radius definition, atomic size generally decreases across the period. Nuclear charge increases, while added electrons enter the same principal shell. Inner-shell shielding stays broadly similar and the additional same-shell shielding does not fully offset the extra nuclear charge. The stronger effective attraction draws the electron cloud closer.
Do not say that shielding is exactly unchanged, or that a new shell is added at every step. Atomic radius is not a sharp physical boundary; values are inferred from particular interactions. Compare data of the same type. A noble-gas van der Waals radius can appear larger than a neighbouring covalent radius because the definitions differ, rather than because the effective nuclear-attraction argument has reversed.
3. Ionic radius: cations, anions and isoelectronic series
A cation is smaller than its parent atom. Na loses its outer 3s electron to form Na⁺ (2,8), removing the n = 3 occupied shell. Mg²⁺ and Al³⁺ also have 10 electrons. With the same electron count and shell arrangement, increasing proton number pulls the cloud closer: Na⁺ > Mg²⁺ > Al³⁺.
An anion is larger than its parent atom. Adding electrons increases electron–electron repulsion and reduces the effective attraction per electron without increasing proton number. P³⁻, S²⁻ and Cl⁻ all have 18 electrons (2,8,8). Their radii decrease in that order as nuclear charge rises. The anions are much larger than the 10-electron cations because the anions retain occupied n = 3 shells.
Do not draw a single smooth decreasing ionic-radius curve through every element. State the ion charges and electron arrangements. Silicon commonly forms covalent compounds; writing oxidation number +4 in SiCl₄ does not prove that it contains isolated Si⁴⁺ ions. Argon does not form an ordinary stable monatomic ion in this introductory comparison.
4. Melting points: explain the change in structure
Na has relatively weak metallic bonding and melts near 98°C. Mg and Al have stronger metallic cohesion, with more delocalised valence electrons per atom and smaller, more highly charged positive cores. This broadly explains their much higher melting points. Mg and Al melt at similar temperatures; crystal structure and other details prevent a simple exact proportionality to electron count.
Silicon has the maximum melting point in this comparison. Its extended tetrahedral covalent network requires substantial disruption of strong covalent bonds to melt. It has no discrete Si molecules separated only by weak London forces.
There is a sharp fall from Si to white P₄. Melting a molecular solid mainly overcomes intermolecular attractions while covalent bonds within molecules remain. S₈ has more electrons and a more polarisable cloud than P₄, producing stronger London attractions and a higher melting point. Molecular shape and packing also contribute.
Cl₂ is a smaller molecular particle with weaker overall dispersion attractions than P₄ or S₈, so its melting point is much lower. Monatomic Ar has particularly weak interatomic London attractions and the lowest melting point here. These comparisons concern the discrete particles and their attractions; counting “covalent bonds per molecule” alone does not explain melting.
| Elemental form | Na | Mg | Al | Si | White P₄ | S₈ | Cl₂ | Ar |
|---|---|---|---|---|---|---|---|---|
| Melting point / °C | 97.8 | 650 | 660.3 | 1414 | 44.2 | 115.2 | −101.5 | −189.3 |
Reference data: Royal Society of Chemistry periodic table. Values are rounded for revision. Allotrope and pressure matter; use supplied examination data for calculations.
5. Electrical conductivity: identify the mobile carriers
Na, Mg and Al conduct as solids and liquids because delocalised electrons can move through the structure and carry charge. More available electrons contribute to the broad trend, but conductivity also depends on their mobility and scattering. Do not claim that Mg must have exactly twice Na’s conductivity or that every neighbouring metal gives a strict increase.
Pure silicon is a semiconductor, not an ordinary metal or a perfect insulator. Most valence electrons participate in covalent bonding, but thermal excitation creates a small population of mobile carriers. Its conductivity increases strongly with temperature and can be changed greatly by doping. A qualitative school graph often places Si far below the metals.
White P₄, S₈, Cl₂ and Ar have no appreciable mobile charged particles under ordinary conditions, so their conductivity is negligible compared with metals. They contain electrons, but those electrons are not freely mobile through the bulk. Black phosphorus has a different extended structure and semiconductor behaviour, so the P₄ qualification matters.
Worked examples
Example 1 · Explain Na versus Cl atomic size
Both atoms occupy n = 3, but Cl has 17 protons compared with Na’s 11. Similar inner shielding means stronger effective attraction across the period, so the cloud contracts. For numerical values compare the same radius definition; do not mix an Na metallic radius and a Cl covalent radius without noting that convention.
Example 2 · Rank Na⁺, Mg²⁺ and Al³⁺
Count electrons: 11 − 1 = 10; 12 − 2 = 10; 13 − 3 = 10. All have arrangement 2,8. Proton numbers increase 11 → 12 → 13, so radius decreases: Na⁺ > Mg²⁺ > Al³⁺. The charges alone are not the complete explanation.
Example 3 · Rank P³⁻, S²⁻ and Cl⁻
Electron counts are 15 + 3 = 18, 16 + 2 = 18 and 17 + 1 = 18. Each has arrangement 2,8,8. Increasing proton number contracts the same electron arrangement, giving P³⁻ > S²⁻ > Cl⁻.
Example 4 · Explain the Si-to-P fall
Si has a giant covalent network; white phosphorus contains P₄ molecules. Melting Si disrupts strong network bonds, whereas melting P₄ mainly overcomes weaker attractions between molecules. The large drop is a change of structure, not evidence that P–P covalent bonds simply vanish.
Example 5 · Read the sulfur–phosphorus difference
Using the rounded data, ΔT = 115.2 − 44.2 = 71.0°C. Explain the higher S₈ value using its larger, more polarisable electron cloud and stronger London attractions, while recognising that packing also matters.
Example 6 · Predict conductivity from structure
A sample has an extended silicon network and very low conductivity at low temperature that rises on heating. This is consistent with a semiconductor: heating increases mobile-carrier population. A metal already has many mobile electrons; typical metallic conductivity decreases on heating because scattering increases.
Common misconceptions and exam tips
“More electrons always means a larger atom.” Across a period the rising effective nuclear attraction dominates.
“All Period 3 ions have the same radius trend.” Separate the 10-electron cations and 18-electron anions.
“Melting P₄ breaks all its covalent bonds.” The main change is between molecules.
“Silicon cannot conduct at all.” It is a semiconductor.
“Any phosphorus sample has the same properties.” Identify the allotrope.
“All eight elements are room-temperature solids.” Chlorine and argon are gases under ordinary room conditions.
For radius questions give electron arrangement, nuclear charge and shielding. For melting state particles, structure, attraction and energy. For conductivity name the mobile charge carrier. Label graph axes and units; do not add measured units to a schematic trend.
Your revision summary
Atomic contraction reflects increasing effective nuclear attraction across n = 3. Cations shrink relative to atoms; anions expand, with separate isoelectronic trends. Melting points depend on the metallic → giant covalent → discrete-particle transition. Metals conduct through delocalised electrons, Si is a semiconductor, and the ordinary right-hand molecular substances and Ar conduct very poorly.
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