Chemical Bonding Revision Guide
GCSE Chemistry | Foundation & Higher
Key Facts
- Atoms bond to achieve a full outer shell of electrons (a stable electron configuration).
- Ionic bonding: Transfer of electrons between a metal and a non-metal.
- Covalent bonding: Sharing of electron pairs between non-metals.
- Metallic bonding: A lattice of positive ions surrounded by a sea of delocalised electrons.
- The type of bonding determines the properties of a substance.
Ionic Bonding
Ionic bonding occurs between metals and non-metals. The metal atom loses electrons to become a positively charged ion (cation). The non-metal atom gains electrons to become a negatively charged ion (anion). The opposite charges attract strongly, forming an ionic bond.
Example: Sodium Chloride (NaCl)
Sodium (Na) has 1 electron in its outer shell. Chlorine (Cl) has 7. Sodium transfers its outer electron to chlorine. Sodium becomes Na+ (2,8 configuration). Chlorine becomes Cl- (2,8,8 configuration). The electrostatic attraction between Na+ and Cl- ions forms a giant ionic lattice.
Properties of Ionic Compounds
- High melting and boiling points: Strong electrostatic forces between ions require lots of energy to break.
- Conduct electricity when dissolved or molten: The ions are free to move and carry charge.
- Do not conduct electricity when solid: The ions are held in fixed positions in the lattice.
- Usually soluble in water.
- Brittle: If layers shift, ions with the same charge line up and repel, causing the crystal to shatter.
Covalent Bonding
Covalent bonding occurs between non-metal atoms. The atoms share one or more pairs of electrons so that each atom achieves a full outer shell. A single bond shares one pair of electrons. A double bond shares two pairs.
Common Examples
- Water (H₂O): Oxygen shares one electron pair with each hydrogen atom (two single bonds).
- Oxygen (O₂): Two oxygen atoms share two electron pairs (a double bond, O=O).
- Methane (CH₄): Carbon shares one electron pair with each of four hydrogen atoms.
- Carbon dioxide (CO₂): Carbon shares two electron pairs with each oxygen (two double bonds, O=C=O).
Simple Molecular Substances
Small covalent molecules (like water, oxygen, methane) have strong covalent bonds within the molecule but weak intermolecular forces between molecules. This gives them:
- Low melting and boiling points: Only weak intermolecular forces need to be overcome (not the covalent bonds).
- Do not conduct electricity: No free electrons or ions to carry charge.
- Often gases or liquids at room temperature.
Giant Covalent Structures
Some covalent substances form giant structures where billions of atoms are bonded together:
- Diamond: Each carbon atom bonds to four others in a rigid tetrahedral structure. Very hard. Very high melting point. Does not conduct electricity (no free electrons).
- Graphite: Each carbon atom bonds to three others in flat layers. The fourth electron is delocalised, so graphite conducts electricity. Layers slide over each other, making graphite soft and slippery (used as a lubricant).
- Silicon dioxide (sand): Giant covalent structure with very high melting point.
Metallic Bonding
In metals, atoms are arranged in a regular lattice. The outer electrons are delocalised (free to move), forming a "sea of electrons" that surrounds positive metal ions. The strong electrostatic attraction between the positive ions and the negative sea of electrons holds the structure together.
Properties of Metals
- Good conductors of electricity: Delocalised electrons are free to move and carry charge.
- Good conductors of heat: Delocalised electrons transfer kinetic energy quickly.
- High melting and boiling points: Strong metallic bonds require lots of energy to break.
- Malleable and ductile: Layers of ions can slide over each other without breaking the bond (unlike ionic compounds, which shatter).
Worked Examples
Example 1: Explain why sodium chloride has a high melting point.
Sodium chloride has a giant ionic lattice structure. There are strong electrostatic forces of attraction between the positive Na+ ions and the negative Cl- ions. A lot of energy is needed to overcome these strong ionic bonds, so the melting point is high.
Example 2: Explain why graphite conducts electricity but diamond does not.
In graphite, each carbon atom bonds to three others, leaving one electron delocalised (free to move). These delocalised electrons can carry charge through the structure. In diamond, each carbon bonds to four others, and all electrons are used in bonding, so there are no free electrons to carry charge.
Example 3: Explain why metals are good conductors of electricity.
Metals have a structure of positive ions surrounded by a sea of delocalised electrons. These delocalised electrons are free to move through the metal lattice and carry electrical charge.
Common Mistakes
- Saying ionic compounds "share" electrons. Ionic bonding involves transfer, not sharing.
- Saying covalent bonds break when a substance melts. In simple molecules, it's the intermolecular forces that break, not the covalent bonds.
- Saying metals conduct because of ions moving. It is the delocalised electrons that carry charge.
- Confusing properties of simple molecular substances with giant covalent structures. They are very different.
- Forgetting that ionic compounds only conduct when dissolved or molten, not when solid.
Exam Tips
- Be able to draw dot and cross diagrams for ionic and covalent compounds. Show only the outer shell electrons.
- When explaining properties, always link the structure to the property. For example: "Ionic compounds have high melting points BECAUSE of the strong electrostatic forces between ions."
- Know the differences between the three types of bonding and be able to compare them in a table.
- Remember: diamond and graphite are both carbon, but they have completely different properties because of their different structures.
- For 6-mark questions, describe the bonding, the structure, and then explain how this gives rise to the properties asked about.
Practice Questions
- Draw a dot and cross diagram for magnesium oxide (MgO).
- Explain why sodium chloride conducts electricity when dissolved in water but not when solid.
- Explain why water has a low boiling point compared to sodium chloride.
- Compare the structures and properties of diamond and graphite.
- Explain why metals are malleable but ionic compounds are brittle.
- A substance has a high melting point, does not conduct electricity when solid, but does conduct when molten. What type of bonding does it have? Explain your answer.
Recommended Revision Guides
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- CGP GCSE Chemistry Complete Revision and Practice — Comprehensive coverage of all topics.
- Hodder GCSE AQA Chemistry Revision Guide — Exam-focused with clear explanations.
- Oxford Revise AQA GCSE Chemistry — Visual revision with exam practice.