Describe the giant covalent structure of silicon(IV) oxide, $mathrm{SiO}_2$

Atoms, Elements and Compounds – Giant Covalent (Network) Solids

Objective

To describe the giant covalent (network) structures of silicon(IV) oxide (SiO₂), diamond and graphite, and to relate these structures to their characteristic physical properties as required by the Cambridge IGCSE Chemistry (0620) syllabus.

1. What is a Giant Covalent (Network) Solid?

• Atoms are linked by strong covalent bonds that extend continuously in three dimensions (or, for graphite, in two‑dimensional layers).
• The entire crystal behaves as one giant molecule; there are no discrete molecules.
• Typical properties: very high melting point, great hardness, brittleness (except graphite), electrical insulation (diamond, SiO₂) or conductivity (graphite), low solubility, and transparency (SiO₂, diamond).

2. Silicon(IV) Oxide – SiO₂

2.1 Structural Description

• Each silicon atom is tetrahedrally coordinated to four oxygen atoms.
• Each oxygen atom is two‑coordinate, acting as a bridge between two silicon atoms (Si–O–Si).
• The basic unit is a SiO₄ tetrahedron; tetrahedra share **corner** oxygen atoms to give an endless 3‑D framework.

2.2 Bonding Details

• Si–O bonds are strong covalent bonds (bond energy ≈ 452 kJ mol⁻¹) with considerable polarity (Si δ⁺, O δ⁻) but no free ions.
• Each Si atom forms four σ‑bonds; each O atom forms two σ‑bonds.
• The Si–O–Si bond angle in quartz is about 144°, giving the network a relatively open structure.

2.3 Quantitative Data

2.4 Physical Properties Linked to Structure

1. Very high melting point (≈ 1710 °C) – breaking a large number of strong Si–O covalent bonds.
2. Hard and brittle – the rigid 3‑D network resists deformation; fracture occurs before bonds can slip.
3. Electrical insulator – electrons are localised in covalent bonds; no free charge carriers.
4. Transparent to visible light – a wide band gap (~9 eV) prevents absorption of visible photons.
5. Low solubility in water – disruption of the extensive Si–O network requires more energy than water can provide.

3. Diamond (Carbon)

3.1 Structural Description

• Each carbon atom is tetrahedrally bonded to four other carbon atoms.
• The tetrahedra share **all** corners, producing a 3‑D network identical in geometry to the SiO₂ framework but without oxygen.

3.2 Bonding Details

• C–C single bonds are covalent with a bond energy of about 347 kJ mol⁻¹.
• All four bonds are equivalent (sp³ hybridisation).

3.3 Physical Properties and Structural Origin

1. Highest known hardness – each carbon is locked in a rigid 3‑D lattice.
2. Very high melting point (≈ 3550 °C) – many strong C–C bonds must be broken.
3. Electrical insulator – no free electrons; large band gap (~5.5 eV).
4. Transparent and high refractive index – wide band gap allows transmission of visible light.

4. Graphite (Carbon)

4.1 Structural Description

• Each carbon atom is bonded to three others in a planar hexagonal sheet (sp² hybridisation).
• Sheets are stacked; adjacent sheets are held together by weak van der Waals forces.
• Within a sheet each carbon forms three σ‑bonds and retains one delocalised π‑electron.

4.2 Bonding Details

• In‑plane C–C σ‑bond energy ≈ 347 kJ mol⁻¹ (as in diamond).
• Delocalised π‑electrons give each layer metallic‑like conductivity.

4.3 Physical Properties and Structural Origin

1. Good electrical conductor (parallel to sheets) – mobile π‑electrons.
2. Soft and lubricating – weak interlayer van der Waals forces allow layers to slide.
3. High melting point (≈ 3600 °C) for the sheets – strong covalent bonds within each sheet.
4. Opaque, black appearance – broad absorption due to π‑electron transitions.

5. Comparison of Giant Covalent Solids

6. Summary

Giant covalent solids are characterised by extensive networks of strong covalent bonds. In SiO₂ the network consists of corner‑sharing SiO₄ tetrahedra, giving rise to a very high melting point, hardness, brittleness, insulating behaviour, transparency and low solubility. Diamond shares the same tetrahedral geometry but with C–C bonds, resulting in extreme hardness and a high melting point. Graphite, by contrast, forms planar sheets of sp²‑bonded carbon; the weak forces between sheets make it soft and a good conductor along the layers. Understanding these structural differences enables students to predict and explain the physical properties required by the Cambridge IGCSE Chemistry syllabus.