Welcome to Structure in Chemistry!
Ever wondered why diamond is one of the hardest substances known to science, while graphite (which is also made entirely of carbon) is so soft that you can use it to write on paper? Or why a copper pipe can be bent into shape without snapping, while a block of salt shatters instantly when hit with a hammer?
The answer lies in one simple rule: Structure dictates properties! In this chapter, we will explore the four major types of crystal structures that you need to master for CCEA AS Chemistry. Don't worry if this feels a bit overwhelming at first — we will break down each structure step-by-step with simple analogies and clear explanations.
The Golden Rule of Physical Properties:
Whenever an exam question asks why a substance has a high melting point, conducts electricity, or is brittle, you must always look at two things:
1. What particles make up the structure? (Ions, atoms, or molecules?)
2. What forces or bonds hold them together? (Ionic, covalent, metallic, or intermolecular forces?)
1. Giant Ionic Lattices
Ionic compounds do not exist as tiny, isolated pairs of ions. Instead, millions of oppositely charged ions pack tightly together in a regular, repeating three-dimensional pattern known as a giant ionic lattice.
Structure of Sodium Chloride (\(\text{NaCl}\))
A classic example you need to know is table salt, \(\text{NaCl}\):
• Each positive sodium ion (\(\text{Na}^+\)) is surrounded octahedrally by \(6\) negative chloride ions (\(\text{Cl}^-\)).
• Each negative chloride ion (\(\text{Cl}^-\)) is surrounded by \(6\) positive sodium ions (\(\text{Na}^+\)).
• This is described as a \(6:6\) co-ordination structure.
Properties of Giant Ionic Structures
1. High Melting and Boiling Points:
• Why? There are very strong electrostatic attractions between oppositely charged ions acting in all directions throughout the giant lattice.
• Explanation: A massive amount of heat energy is required to overcome these strong ionic bonds.
• Comparison: Magnesium oxide (\(\text{MgO}\)) has a much higher melting point (\(\approx 2852\text{ }^\circ\text{C}\)) than sodium chloride (\(\text{NaCl}\), \(\approx 801\text{ }^\circ\text{C}\)) because \(\text{Mg}^{2+}\) and \(\text{O}^{2-}\) ions have higher charges than \(\text{Na}^+\) and \(\text{Cl}^-\), leading to stronger electrostatic attraction.
2. Electrical Conductivity:
• As a solid: Does not conduct electricity because the ions are locked tightly in fixed positions within the lattice and are not free to move.
• When molten (liquid) or dissolved in water (aqueous): Conducts electricity because the ionic lattice breaks down, allowing the ions to become mobile charge carriers.
3. Brittleness:
• Why? When a mechanical force or stress is applied to an ionic crystal, layers of ions slide over each other.
• Ions of the same charge are forced next to each other (e.g., \(\text{Na}^+\) next to \(\text{Na}^+\)).
• The resulting strong electrostatic repulsion shatters the crystal apart.
4. Solubility in Water:
• Most ionic compounds dissolve in polar solvents like water.
• The partially negative oxygen atoms of \(\text{H}_2\text{O}\) attract positive cations (\(\text{Na}^+\)), while the partially positive hydrogen atoms attract negative anions (\(\text{Cl}^-\)). This process releases hydration energy which helps break apart the lattice.
Key Takeaway for Ionic Lattices:
High melting points, brittle, conduct electricity ONLY when molten or in solution because ions must be free to move.
2. Giant Metallic Lattices
Metals have a unique structure that gives them their characteristic shine, strength, and flexibility.
Structure of a Metal
A metallic lattice consists of a regular, repeating arrangement of positive metal cations surrounded by a 'sea' of delocalised electrons.
• Delocalised electrons are outer-shell electrons that are no longer tied to any single atom and can move freely throughout the entire structure.
• Metallic bonding is the strong electrostatic attraction between the positive metal ions and the delocalised electrons.
Properties of Giant Metallic Structures
1. High Melting and Boiling Points:
• The electrostatic attraction between the positive metal cations and the sea of delocalised electrons is very strong and requires a large amount of energy to break.
• Trend: Melting point increases with higher ion charge and smaller ionic radius because the delocalised electron density is greater. For example, aluminium (\(\text{Al}^{3+}\)) has a higher melting point than sodium (\(\text{Na}^+\)).
2. Excellent Electrical and Thermal Conductivity:
• Metals conduct electricity in both solid and molten states.
• Why? The delocalised electrons are free to move and carry electrical charge or transfer kinetic thermal energy rapidly through the lattice.
3. Malleability and Ductility:
• Malleable: Can be hammered or rolled into thin sheets.
• Ductile: Can be drawn out into long wires.
• Why? The layers of positive metal cations can easily slide over one another when a force is applied without breaking the metallic bond, because the sea of delocalised electrons moves with them and maintains the electrostatic attraction.
Key Takeaway for Metallic Lattices:
Always mention the sea of delocalised electrons when explaining electrical conductivity or malleability in metals!
3. Giant Covalent (Macromolecular) Structures
In giant covalent structures, hundreds of thousands of atoms are linked together in a continuous network by strong covalent bonds.
A. Diamond
• Structure: Each carbon atom is covalently bonded to four other carbon atoms in a rigid, three-dimensional tetrahedral arrangement. The bond angles are \(109.5^\circ\).
• Hardness: Extremely hard because breaking the crystal requires snapping many strong covalent bonds.
• Melting Point: Exceptionally high (over \(3500\text{ }^\circ\text{C}\)) due to the vast network of strong covalent bonds.
• Electrical Conductivity: Non-conductor because all four outer-shell valence electrons of each carbon are fixed in covalent bonds; there are no free electrons or ions to carry charge.
B. Graphite
• Structure: Each carbon atom is covalently bonded to three other carbon atoms in flat, hexagonal planar rings, forming 2D layers. The bond angles are \(120^\circ\).
• Weak Forces: The individual layers are held together by weak van der Waals (induced dipole-dipole) forces.
• Soft and Slippery: The weak forces between layers allow the layers to slide easily over each other. This makes graphite an excellent solid lubricant and perfect for pencil leads.
• Electrical Conductivity: Good conductor because each carbon atom has one delocalised electron (the 4th valence electron) that is free to move parallel through the layers.
• High Melting Point: Despite having weak forces between layers, graphite has a huge melting point because the covalent bonds within each layer are extremely strong.
C. Graphene
• Structure: A single, two-dimensional sheet of graphite just one atom thick.
• Properties: Incredibly strong for its weight, transparent, and an outstanding electrical conductor because of freely moving delocalised electrons across the 2D plane.
D. Silicon Dioxide (\(\text{SiO}_2\), Silica / Quartz)
• Structure: A giant macromolecular structure similar to diamond.
• Each silicon (\(\text{Si}\)) atom is covalently bonded to four oxygen atoms in a tetrahedral arrangement.
• Each oxygen (\(\text{O}\)) atom is bonded to two silicon atoms.
• Properties: Very hard, very high melting point, and a non-conductor of electricity.
Key Takeaway for Giant Covalent Structures:
Diamond and quartz do NOT conduct electricity because all valence electrons are fixed in bonds. Graphite and graphene DO conduct because each carbon uses only \(3\) of its \(4\) electrons for bonding, leaving \(1\) delocalised electron per atom!
4. Simple Molecular Structures
Simple molecular substances consist of individual molecules held internally by strong covalent bonds, but the separate molecules are held together in a crystal lattice by relatively weak intermolecular forces.
A. Iodine (\(\text{I}_2\))
• Structure: Diatomic \(\text{I}_2\) molecules arranged in a regular crystal lattice (face-centred cubic).
• Bonding: Strong covalent bonds hold the two iodine atoms together inside each \(\text{I}_2\) molecule, but only weak van der Waals forces hold the separate molecules to one another.
• Properties: Low melting point; readily sublimes (turns directly from a grey-black solid to a purple gas) upon gentle heating because very little energy is needed to overcome the weak intermolecular forces.
• Conductivity: Non-conductor because there are no free-moving ions or delocalised electrons.
B. Ice (Solid Water, \(\text{H}_2\text{O}\))
• Structure: In ice, each water molecule forms hydrogen bonds with four neighbouring molecules, creating an open, hexagonal 3D tetrahedral network.
• Density: The open lattice holds molecules further apart than in liquid water. Therefore, ice is less dense than liquid water, which is why icebergs float!
C. Buckminsterfullerene (\(\text{C}_{60}\))
• Structure: A simple molecular allotrope of carbon where \(60\) carbon atoms form a hollow sphere (like a football) made of \(20\) hexagons and \(12\) pentagons.
• Properties: Has a much lower melting point than diamond or graphite because you only need to overcome weak van der Waals forces between the spherical \(\text{C}_{60}\) cages, not covalent bonds.
Key Takeaway for Simple Molecular Structures:
When melting simple molecular crystals, you are ONLY breaking weak intermolecular forces (van der Waals forces or hydrogen bonds), NEVER the strong internal covalent bonds!
5. Summary & Comparison Guide
Here is a quick reference guide to keep the four structures clear in your mind:
1. Giant Ionic (e.g., \(\text{NaCl}\), \(\text{MgO}\))
• Particles: Positive and negative ions.
• Forces broken on melting: Strong ionic bonds.
• Melting point: High.
• Electrical conductivity: Solid: No | Molten/Aqueous: Yes.
• Other features: Brittle, often water-soluble.
2. Giant Metallic (e.g., \(\text{Mg}\), \(\text{Cu}\), \(\text{Fe}\))
• Particles: Positive metal ions in a sea of delocalised electrons.
• Forces broken on melting: Strong metallic bonds.
• Melting point: High to moderate.
• Electrical conductivity: Solid: Yes | Molten: Yes.
• Other features: Malleable, ductile, shiny.
3. Giant Covalent (e.g., Diamond, Graphite, \(\text{SiO}_2\))
• Particles: Non-metal atoms.
• Forces broken on melting: Strong covalent bonds.
• Melting point: Very high.
• Electrical conductivity: Solid: No (except Graphite/Graphene) | Molten: No.
• Other features: Very hard (except graphite, which is soft/slippery).
4. Simple Molecular (e.g., \(\text{I}_2\), \(\text{H}_2\text{O}\), \(\text{C}_{60}\))
• Particles: Small discrete molecules.
• Forces broken on melting: Weak intermolecular forces.
• Melting point: Low.
• Electrical conductivity: Solid: No | Molten: No.
• Other features: Soft, volatile, low boiling points.
Common Pitfalls to Avoid in CCEA Exams
• Mistake 1: Stating that ionic solids conduct electricity because "electrons can move".
Correction: In ionic compounds, conductivity is due to mobile ions, not electrons!
• Mistake 2: Saying that "covalent bonds break when water boils or iodine melts".
Correction: Only weak intermolecular forces (van der Waals or hydrogen bonds) break during phase changes of simple molecules. Covalent bonds inside the molecule remain intact.
• Mistake 3: Forgetting why graphite conducts electricity.
Correction: Always mention that each carbon is bonded to only three others, leaving one delocalised electron per carbon atom that is free to move along the layers.