Across Period 3 (from sodium to argon), elements exhibit distinct physical trends. These trends are determined by the electronic structure of the atoms, their nuclear charge, and the type of structure and bonding they form.
🔑 Key Principle
Physical properties across Period 3 are dictated by structure and bonding. The period is divided into three distinct types of structure: giant metallic lattices (Na, Mg, Al), a giant covalent lattice (Si), and simple molecular lattices (P₄, S₈, Cl₂, Ar).
Trend 1: Atomic Radius
The atomic radius is a measure of the size of an atom. Across Period 3, the atomic radius decreases from sodium (\(\text{Na}\)) to chlorine (\(\text{Cl}\)).
Explanation:
- Nuclear Charge Increases: The number of protons in the nucleus increases by one with each successive element across the period (from +11 in \(\text{Na}\) to +17 in \(\text{Cl}\)).
- Shielding remains roughly constant: Electrons are added to the same principal energy level (the 3rd shell). The inner shells of core electrons (\(1\text{s}^2 2\text{s}^2 2\text{p}^6\)) provide a constant level of shielding.
- Stronger Attraction: The increased nuclear charge attracts the outer shell electrons more strongly, pulling them closer to the nucleus and decreasing the atomic radius.
Argon (\(\text{Ar}\)) is typically excluded from covalent radius comparisons. As a noble gas, it exists as monatomic atoms and does not form covalent bonds under standard conditions. Thus, its size is measured via its van der Waals radius, which appears artificially larger because the weak intermolecular forces allow atoms to remain further apart than covalently bonded atoms.
Trend 2: First Ionisation Energy
The first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions. Generally, the first ionisation energy increases across Period 3 from sodium to argon. This is because of the increased nuclear charge, smaller atomic radius, and constant shielding, making it harder to remove outer electrons.
However, there are two key exceptions to this general increase:
- Dip at Aluminium (\(\text{Al}\)): Aluminium's outer electron configuration is \([\text{Ne}]3\text{s}^2 3\text{p}^1\), whereas Magnesium's is \([\text{Ne}]3\text{s}^2\). The outer electron in \(\text{Al}\) is in a 3p sub-shell, which is higher in energy and shielded by the 3s sub-shell. It is therefore easier to remove than a 3s electron in \(\text{Mg}\).
- Dip at Sulfur (\(\text{S}\)): Sulfur's outer configuration is \([\text{Ne}]3\text{s}^2 3\text{p}^4\), whereas Phosphorus's is \([\text{Ne}]3\text{s}^2 3\text{p}^3\). In sulfur, the fourth p electron shares an orbital with another electron. The mutual repulsion between these paired electrons in the same orbital makes the outer electron easier to remove.
Trend 3: Melting Points
Unlike radius and ionisation energy, melting point trends cannot be explained solely by atomic trends. Instead, they depend on the physical structure and bonding type of each substance.
1. Na, Mg, Al: Metallic Structure
Sodium, magnesium, and aluminium are metals with giant metallic lattices. Their melting points increase across this series because metallic bonding gets progressively stronger:
- Charge on Ion Increases: The charge on the metal ions increases (\(\text{Na}^+\) to \(\text{Mg}^{2+}\) to \(\text{Al}^{3+}\)).
- More Delocalised Electrons: The number of delocalised electrons donated per atom increases (1 in \(\text{Na}\), 2 in \(\text{Mg}\), 3 in \(\text{Al}\)).
- Smaller Ionic Radius: The radius of the positive metal ion decreases, allowing the positive nuclei and delocalised electrons to get closer together.
This leads to a stronger electrostatic attraction between positive metal ions and the surrounding sea of delocalised electrons, requiring more energy to overcome.
Metallic bonding is the strong electrostatic attraction between a lattice of positive metal ions and a sea of delocalised valence electrons.
2. Si: Giant Covalent Structure
Silicon is a macromolecular metalloid with a giant covalent structure (similar to diamond). Its melting point is exceptionally high because melting it requires breaking many strong, localized covalent silicon-silicon (\(\text{Si}-\text{Si}\)) bonds. Breaking these strong covalent bonds requires a vast amount of thermal energy.
A giant covalent structure (macromolecular) is a network of atoms linked together by strong covalent bonds extending throughout a giant 3D lattice.
3. P₄, S₈, Cl₂, Ar: Simple Molecular Structures
Phosphorus, sulfur, chlorine, and argon exist as simple molecules or monatomic atoms. Their structures are simple molecular lattices:
- Phosphorus exists as \(\text{P}_4\) molecules.
- Sulfur exists as \(\text{S}_8\) molecules.
- Chlorine exists as \(\text{Cl}_2\) molecules.
- Argon exists as monatomic \(\text{Ar}\) atoms.
When these substances melt, it is not the covalent bonds within the molecules that are broken. Instead, you only need to overcome the weak intermolecular forces (London/van der Waals forces) holding the molecules together. Since these forces are weak, these elements have relatively low melting points.
The trend in melting points among these non-metals is: \(\text{S}_8 > \text{P}_4 > \text{Cl}_2 > \text{Ar}\). This trend is explained by molecule size:
- Sulfur (\(\text{S}_8\)) is the largest molecule with the most electrons (128 electrons), resulting in the strongest van der Waals forces.
- Phosphorus (\(\text{P}_4\)) has 60 electrons, so its intermolecular forces are weaker than sulfur's.
- Chlorine (\(\text{Cl}_2\)) has 34 electrons, so its intermolecular forces are weaker still.
- Argon (\(\text{Ar}\)) is monatomic with only 18 electrons, resulting in the weakest intermolecular forces of all.
Always state the structure and bonding type first before explaining trends. In exam questions, never state that "covalent bonds are broken" when melting sulfur, phosphorus, or chlorine. You are only overcoming weak intermolecular (van der Waals) forces. Silicon is the only non-metal in Period 3 where covalent bonds actually break during melting.
Period 3 Physical Properties Summary
| Element | Atomic Number | Atomic Radius / pm | 1st Ionisation Energy / kJ mol⁻¹ | Melting Point / K | Structure type | Primary Bonding Type |
|---|---|---|---|---|---|---|
| Sodium (Na) | 11 | 186 | 496 | 371 | Giant metallic | Metallic |
| Magnesium (Mg) | 12 | 160 | 738 | 923 | Giant metallic | Metallic |
| Aluminium (Al) | 13 | 143 | 578 | 933 | Giant metallic | Metallic |
| Silicon (Si) | 14 | 118 | 789 | 1687 | Giant covalent | Covalent |
| Phosphorus (P₄) | 15 | 110 | 1012 | 317 | Simple molecular | Covalent (intramolecular), van der Waals (intermolecular) |
| Sulfur (S₈) | 16 | 102 | 1000 | 392 | Simple molecular | Covalent (intramolecular), van der Waals (intermolecular) |
| Chlorine (Cl₂) | 17 | 99 | 1251 | 172 | Simple molecular | Covalent (intramolecular), van der Waals (intermolecular) |
| Argon (Ar) | 18 | (188 van der Waals) | 1521 | 84 | Monatomic | van der Waals (intermolecular) |
- Explain why aluminium has a higher melting point than sodium.
- Explain why sulfur has a higher melting point than phosphorus.
Part 1: Aluminium vs Sodium
- Both aluminium and sodium form giant metallic lattices with metallic bonding.
- The aluminium ion (\(\text{Al}^{3+}\)) has a higher charge than the sodium ion (\(\text{Na}^+\)).
- Aluminium contributes more delocalised electrons (3 per atom) than sodium (1 per atom).
- The \(\text{Al}^{3+}\) ion is smaller than the \(\text{Na}^+\) ion.
- Therefore, the electrostatic attraction between positive ions and delocalised electrons is stronger in aluminium, requiring more thermal energy to overcome.
Part 2: Sulfur vs Phosphorus
- Both sulfur and phosphorus exist as simple molecular lattices with weak van der Waals forces between molecules.
- Sulfur exists as larger \(\text{S}_8\) molecules, whereas phosphorus exists as smaller \(\text{P}_4\) molecules.
- Because \(\text{S}_8\) molecules contain more electrons (128 electrons) than \(\text{P}_4\) molecules (60 electrons), they form stronger instantaneous dipole-induced dipole (van der Waals) forces.
- More thermal energy is required to overcome the stronger intermolecular forces in sulfur, resulting in a higher melting point.
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