Bonding, Structure, and Properties
Ionic, covalent, and metallic bonding; giant and simple molecular structures.
In this topic
Ionic Bonding
Ionic bonding occurs between metals and non-metals. Metal atoms lose electrons to become positive ions (cations). Non-metal atoms gain electrons to become negative ions (anions).
The oppositely charged ions are attracted to each other by strong electrostatic forces — this is the ionic bond.
Ionic compounds form giant ionic lattices — a regular arrangement of alternating positive and negative ions in all directions.
Properties of ionic compounds: - High melting and boiling points (strong bonds) - Conduct electricity when molten or dissolved (ions can move) - Do NOT conduct electricity when solid (ions are fixed) - Usually soluble in water
Key Points
- Metal → positive ion (loses electrons)
- Non-metal → negative ion (gains electrons)
- Ionic compounds are held together by electrostatic attraction
- Dot and cross diagrams show electron transfer
Example Questions
3Explain why sodium chloride has a high melting point.
Sodium chloride has a giant ionic lattice structure with strong electrostatic forces of attraction between the oppositely charged Na⁺ and Cl⁻ ions in all directions. A large amount of energy is needed to overcome these strong ionic bonds.
[3 marks]
2Explain why solid sodium chloride does not conduct electricity, but molten sodium chloride does.
In solid NaCl, the ions are held in fixed positions in the lattice and cannot move to carry charge. When molten, the ions are free to move and can carry charge, allowing electrical conduction.
[2 marks]
Covalent Bonding
Covalent bonding occurs between non-metal atoms. Atoms share pairs of electrons to achieve full outer shells.
A single covalent bond = one shared pair of electrons A double covalent bond = two shared pairs of electrons
Examples: - H₂: H—H (single bond) - O₂: O=O (double bond) - H₂O: two O—H single bonds - CO₂: two C=O double bonds
Simple molecular structures (like H₂O, CO₂): - Low melting/boiling points (weak intermolecular forces between molecules) - Do not conduct electricity (no free electrons or ions) - The covalent bonds within molecules are strong, but the forces between molecules are weak
Key Points
- Covalent = sharing electrons between non-metals
- It's the INTERMOLECULAR forces that are weak, not the covalent bonds
- More electrons shared = stronger bond (double > single)
- Giant covalent structures (diamond, graphite, silicon dioxide) have very high melting points
Example Questions
3Explain why water (H₂O) has a low boiling point compared to sodium chloride.
Water is a simple molecular compound. The covalent bonds within each molecule are strong, but the intermolecular forces (forces between molecules) are weak. Only a small amount of energy is needed to overcome these weak intermolecular forces. NaCl has strong ionic bonds throughout its giant lattice structure, requiring much more energy to break.
[3 marks]
3Explain why diamond has a very high melting point.
Diamond has a giant covalent structure where each carbon atom is covalently bonded to four other carbon atoms in a rigid 3D tetrahedral arrangement. There are many strong covalent bonds throughout the structure, and a very large amount of energy is needed to break them all.
[3 marks]
Metallic Bonding
In metals, atoms are arranged in a regular pattern (lattice). The outer electrons are delocalised — they are free to move through the whole structure, forming a 'sea of electrons'.
Metallic bonding is the strong electrostatic attraction between the positive metal ions and the sea of delocalised electrons.
Properties of metals: - High melting and boiling points (strong metallic bonds) - Good conductors of electricity (delocalised electrons carry charge) - Good conductors of heat (delocalised electrons transfer energy) - Malleable and ductile (layers of ions can slide over each other)
Alloys are mixtures of metals (or metal with non-metal). They are harder than pure metals because the different-sized atoms disrupt the regular layers, preventing them from sliding.
Key Points
- Delocalised electrons = free to move = conduct electricity
- Pure metals are soft — regular layers can slide
- Alloys are harder — irregular atoms disrupt sliding
- Examples of alloys: steel (iron + carbon), brass (copper + zinc), bronze (copper + tin)
Example Questions
2Explain why metals are good conductors of electricity.
Metals have delocalised electrons that are free to move through the structure. When a potential difference is applied, these electrons can flow, carrying charge through the metal and forming an electric current.
[2 marks]
2Explain why alloys are harder than pure metals.
In a pure metal, the atoms are arranged in regular layers that can easily slide over each other. In an alloy, atoms of different sizes are mixed in, which disrupts the regular arrangement and prevents the layers from sliding easily, making the alloy harder.
[2 marks]