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This lesson covers covalent bonding as required by the AQA GCSE Combined Science Trilogy specification (8464), section 4.2.2. You need to understand how covalent bonds form, draw dot-and-cross diagrams for simple molecules, and recognise single and double covalent bonds.
A covalent bond is a shared pair of electrons between two atoms. Covalent bonding occurs between non-metal atoms. Each atom contributes one electron to the shared pair, and the shared electrons are attracted to the nuclei of both atoms, holding them together.
Covalent bonding happens because non-metal atoms need to gain electrons to achieve a full outer shell, but neither atom is willing to give up electrons completely (as metals do in ionic bonding). Instead, they share electron pairs.
Exam Tip (AQA 8464): The definition of a covalent bond is "a shared pair of electrons between two atoms." Make sure you say "shared pair" — a single electron is not a covalent bond.
Each atom in a covalent bond contributes one electron to the shared pair (in a standard covalent bond). The shared pair of electrons is attracted to the positive nuclei of both atoms, which holds the atoms together.
The number of covalent bonds an atom forms depends on how many electrons it needs to achieve a full outer shell:
| Element | Group | Electrons in Outer Shell | Electrons Needed | Bonds Formed |
|---|---|---|---|---|
| Hydrogen (H) | 1 | 1 | 1 | 1 |
| Carbon (C) | 4 | 4 | 4 | 4 |
| Nitrogen (N) | 5 | 5 | 3 | 3 |
| Oxygen (O) | 6 | 6 | 2 | 2 |
| Fluorine (F) | 7 | 7 | 1 | 1 |
| Chlorine (Cl) | 7 | 7 | 1 | 1 |
H⋅+⋅H→H–H
Some atoms need to share two pairs of electrons to achieve a full outer shell. This forms a double covalent bond, shown as two lines (=) in a displayed formula.
O=O
graph TD
A["Types of Covalent Bond"] --> B["Single bond<br/>One shared pair<br/>e.g. H–H, H–Cl"]
A --> C["Double bond<br/>Two shared pairs<br/>e.g. O=O, O=C=O"]
style A fill:#2c3e50,color:#fff
style B fill:#3498db,color:#fff
style C fill:#e67e22,color:#fff
For the AQA 8464 exam, you need to draw dot-and-cross diagrams for covalent molecules. The rules are:
| Molecule | Formula | Bond Type | Lone Pairs |
|---|---|---|---|
| Hydrogen | H₂ | Single | None |
| Water | H₂O | 2 × Single | 2 on O |
| Ammonia | NH₃ | 3 × Single | 1 on N |
| Methane | CH₄ | 4 × Single | None |
| Hydrogen chloride | HCl | Single | 3 on Cl |
| Oxygen | O₂ | Double | 2 on each O |
| Carbon dioxide | CO₂ | 2 × Double | 2 on each O |
| Nitrogen | N₂ | Triple | 1 on each N |
Nitrogen has 5 outer electrons and needs 3 more. Two nitrogen atoms share three pairs of electrons, forming a triple bond (N≡N). Each nitrogen has 1 lone pair.
Exam Tip: When drawing dot-and-cross diagrams, count the total electrons around each atom to check your answer. Hydrogen should have 2, and all other non-metals listed above should have 8 (the octet rule).
graph TD
A["How to identify<br/>the bonding type"] --> B{"Metal + Non-metal?"}
B -->|"Yes"| C["IONIC BONDING<br/>Transfer of electrons"]
B -->|"No"| D{"Two non-metals?"}
D -->|"Yes"| E["COVALENT BONDING<br/>Sharing of electrons"]
D -->|"No — two metals"| F["METALLIC BONDING<br/>Delocalised electrons"]
style A fill:#2c3e50,color:#fff
style C fill:#e74c3c,color:#fff
style E fill:#3498db,color:#fff
style F fill:#f39c12,color:#fff
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