The physical properties of the Period 3 oxides, such as melting point and electrical conductivity, are determined directly by their structure and bonding. As we move across Period 3, the type of bonding transitions from ionic on the left, to giant covalent in the middle, and finally to simple molecular on the right. This change does not follow a simple linear trend but exhibits a dramatic drop in melting point when transitioning from macromolecular to simple molecular structures.
Key Principle: Structure determines physical properties
The high melting points of Na₂O, MgO, and Al₂O₃ are due to their giant ionic structures. The high melting point of SiO₂ is due to its giant covalent structure. The low melting points of P₄O₁₀, SO₂, and SO₃ are due to their simple molecular structures held together by weak intermolecular forces.
1. Giant Ionic Oxides: Na₂O and MgO
Sodium oxide and magnesium oxide exist as giant ionic lattices. These structures are held together by strong electrostatic attractions between oppositely charged ions, which require large amounts of energy to overcome.
- Sodium Oxide (Na₂O): Contains Na⁺ and O²⁻ ions in a giant ionic lattice. It has a high melting point (1405 K).
- Magnesium Oxide (MgO): Contains Mg²⁺ and O²⁻ ions in a giant ionic lattice. It has a very high melting point (3125 K).
Why does MgO have a much higher melting point than Na₂O?
The difference in melting point is due to the strength of the ionic bonds, which is governed by lattice enthalpy. This is explained by two factors:
- Ionic Charge: The Mg²⁺ ion has a higher charge (+2) than the Na⁺ ion (+1). Therefore, the electrostatic attraction between Mg²⁺ and O²⁻ is much stronger than between Na⁺ and O²⁻.
- Ionic Radius: The Mg²⁺ ion has a smaller ionic radius than the Na⁺ ion because it has lost more electrons and has a higher proton-to-electron ratio (greater nuclear pull). This allows the ions to pack closer together, further strengthening the electrostatic forces.
When comparing the melting points of ionic compounds, you must always mention both the charge of the cations (e.g. Mg²⁺ is +2, Na⁺ is +1) and their size (ionic radius). Do not just state that one has stronger ionic bonds without explaining why using these two factors.
2. Aluminium Oxide (Al₂O₃): Ionic with Covalent Character
Aluminium oxide has a giant structure that is intermediate between ionic and covalent. Because the Al³⁺ ion is small and carries a high +3 charge, it has a very high charge density. Consequently, it is highly polarising.
Charge density is a measure of the concentration of electric charge on an ion, calculated as the ratio of the ion's charge to its volume (or ionic radius).
The distortion of the electron cloud of an anion by a nearby cation. A cation with a high charge density (small radius and high charge) is highly polarising and pulls the valence electrons of the anion towards itself, introducing covalent character into the ionic bond.
The Al³⁺ cation polarises the O²⁻ anion, distorting its electron cloud. This causes electron density to build up between the nuclei, giving the bonds significant covalent character. Because of this, the ionic lattice of Al₂O₃ is slightly less stable than expected for a purely ionic compound of these charges, contributing to its melting point (2345 K) being lower than that of MgO, despite the higher charge on the aluminium ion.
3. Silicon Dioxide (SiO₂): Giant Covalent
Silicon dioxide (silica) exists as a giant covalent (macromolecular) structure. Each silicon atom is covalently bonded to four oxygen atoms in a tetrahedral arrangement, and each oxygen atom is bonded to two silicon atoms. To melt silicon dioxide, many strong Si-O covalent bonds throughout the giant structure must be broken. This requires a large amount of energy, resulting in a high melting point (1883 K).
Do not confuse giant covalent structures with simple molecular structures. Students often write that silicon dioxide has a high melting point because of "strong intermolecular forces". There are no molecules and thus no intermolecular forces in silicon dioxide. It is a single giant network held together solely by covalent bonds.
4. Simple Molecular Oxides: P₄O₁₀, SO₂, and SO₃
The oxides of phosphorus and sulfur are non-metal oxides with simple molecular structures. They consist of discrete molecules held together by weak intermolecular forces:
- Phosphorus(V) Oxide (P₄O₁₀): A relatively large molecule consisting of four phosphorus and ten oxygen atoms. Because it is a large molecule with many electrons, it forms relatively strong London forces between molecules, giving it a melting point of 573 K (sublimes).
- Sulfur Dioxide (SO₂) and Sulfur Trioxide (SO₃): Small gaseous or volatile liquid molecules. They have very weak intermolecular forces (London forces and permanent dipole-dipole attractions in SO₂). Consequently, they have very low melting points (SO₂: 200 K; SO₃: 290 K).
5. Electrical Conductivity
The electrical conductivity of these oxides depends on the presence of mobile charged particles (ions or delocalised electrons):
- Ionic Oxides (Na₂O, MgO, Al₂O₃): Do not conduct electricity in the solid state because the ions are fixed in position within the giant lattice. However, they do conduct electricity when molten (liquid) because the electrostatic attractions are broken, leaving the ions free to move and carry charge.
- Giant Covalent Oxide (SiO₂): Does not conduct electricity in either solid or molten states. There are no ions, and all valence electrons are tightly held in localised covalent bonds.
- Simple Molecular Oxides (P₄O₁₀, SO₂, SO₃): Do not conduct electricity in any state because they consist of neutral molecules with no free ions or delocalised electrons.
6. Summary of Structure, Bonding, and Properties
| Oxide | Melting Point (K) | Structure type | Bonding type | Conductivity (Solid) | Conductivity (Liquid) |
|---|---|---|---|---|---|
| Na₂O | 1405 | Giant ionic lattice | Ionic | None | Good |
| MgO | 3125 | Giant ionic lattice | Ionic | None | Good |
| Al₂O₃ | 2345 | Giant ionic lattice | Ionic (covalent character) | None | Good |
| SiO₂ | 1883 | Giant covalent | Covalent | None | None |
| P₄O₁₀ | 573 | Simple molecular | Covalent (intermolecular forces) | None | None |
| SO₂ | 200 | Simple molecular | Covalent (intermolecular forces) | None | None |
| SO₃ | 290 | Simple molecular | Covalent (intermolecular forces) | None | None |
Answer:
- Silicon dioxide has a giant covalent structure. To melt it, many strong covalent bonds between silicon and oxygen atoms must be broken, which requires a large amount of thermal energy.
- Sulfur trioxide has a simple molecular structure. To melt it, only the weak intermolecular forces (London forces) between the sulfur trioxide molecules must be overcome. This requires very little thermal energy.
Answer:
- Both compounds exist as giant ionic lattices.
- Magnesium ions (Mg²⁺) have a higher charge (+2) than sodium ions (Na⁺, +1).
- Magnesium ions are also smaller (have a smaller ionic radius) than sodium ions.
- This results in much stronger electrostatic attractions between the oppositely charged ions (Mg²⁺ and O²⁻) in the magnesium oxide lattice compared to the sodium oxide lattice, requiring significantly more energy to break.
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