3.5 Covalent Bonding

A free GCSE Chemistry lesson for AQA (8462) from OBA School: shared pairs of electrons, dot-and-cross diagrams for simple molecules, and single, double and triple bonds. Watch the video, read the explanation and worked examples, then try the revision questions. You can work through the full lesson, with every video, practice quizzes and printable worksheets with mark schemes, in our free online course.

What you will learn

• Describe a covalent bond as a shared pair of electrons between non-metal atoms
• Draw dot-and-cross diagrams for H₂, Cl₂, HCl, H₂O, NH₃, CH₄, O₂, N₂ and CO₂
• Identify single, double and triple covalent bonds
• Represent molecules with displayed formulae and 3D models, and give the limitations of each
• Work out a molecular formula from a model or diagram

Key words

• Covalent bond: A shared pair of electrons between two atoms. It is a strong bond.
• Molecule: A particle made of two or more atoms held together by covalent bonds.
• Single / double / triple bond: One, two or three shared pairs of electrons between the same two atoms.
• Lone pair: A pair of outer-shell electrons that is not shared in a bond.
• Displayed formula: A diagram showing every atom and every bond, with each bond drawn as a line.
• Molecular formula: The actual number of atoms of each element in one molecule, e.g. CH₄.

Sharing electrons

When non-metal atoms bond with each other, they share pairs of electrons. Each shared pair is a covalent bond. By sharing, each atom gets a full outer shell (2 for hydrogen, 8 for most others).

Key fact: A covalent bond is a shared pair of electrons between two atoms. Covalent bonds are strong. The positive nuclei of both atoms attract the shared electrons.

Covalently bonded substances can be small molecules (e.g. water), very large molecules such as polymers (e.g. poly(ethene)), or giant covalent structures (e.g. diamond and silicon dioxide).

• H (Group 1); Outer electrons 1; Electrons needed 1; Covalent bonds usually formed 1
• Cl (Group 7); Outer electrons 7; Electrons needed 1; Covalent bonds usually formed 1
• O (Group 6); Outer electrons 6; Electrons needed 2; Covalent bonds usually formed 2
• N (Group 5); Outer electrons 5; Electrons needed 3; Covalent bonds usually formed 3
• C (Group 4); Outer electrons 4; Electrons needed 4; Covalent bonds usually formed 4

Exam tip: Ionic or covalent? Look at the elements. Metal + non-metal (e.g. lithium oxide, iron(II) chloride) → ionic. Non-metals only (e.g. hydrogen iodide, sulfur dioxide, nitrogen trichloride) → covalent.

Molecules with single bonds

In a dot-and-cross diagram, the outer shells overlap. Shared electrons go in the overlap. Electrons from one atom are dots and from the other are crosses.

Worked example: checking a dot-and-cross diagram

In ammonia, count the electrons around nitrogen: 3 shared pairs + 1 lone pair = 8 ✓
Each hydrogen has 1 shared pair = 2 electrons ✓ (full first shell).

Answer: all atoms have full outer shells, so the diagram is correct

By sharing, each atom reaches the electronic structure of a noble gas. In hydrogen chloride, H has 2 outer electrons, like helium, and Cl has 2,8,8, like argon. In methane, C has 2,8, like neon.

Double and triple bonds

Sometimes atoms must share two or three pairs to fill their outer shells. Each oxygen atom needs 2 more electrons, so O₂ has a double bond. Each nitrogen atom needs 3, so N₂ has a triple bond. Carbon dioxide has two double bonds.

Exam tip: Count the shared electrons in the overlap: 2 = single, 4 = double, 6 = triple bond.

Other ways to show molecules

• dot-and-cross; Shows which atom each electron came from; lone pairs; Does not show the shape of the molecule; all electrons are really identical
• displayed formula, e.g. H−O−H, O=C=O, N≡N; Shows every atom and bond, and bond type (−, =, ≡); Does not show the 3D shape; lone pairs
• ball-and-stick (3D); Shows the shape and bond angles; Does not show which electrons are shared; atoms are not really far apart
• molecular formula; Shows the number of each type of atom; Does not show how the atoms are joined

Worked example: molecular formula from a model

A model shows a molecule with 2 carbon atoms, 6 hydrogen atoms and 1 oxygen atom.

Answer: C₂H₆O (ethanol, often written C₂H₅OH)

Common mistakes

• Saying atoms in a covalent bond transfer electrons. They share them.
• Drawing one shared electron for a bond. A covalent bond is a shared pair: one dot and one cross.
• Forgetting the lone pairs on oxygen, nitrogen and chlorine, so the outer shell has fewer than 8 electrons.
• Drawing more than 2 electrons around a hydrogen atom.
• Saying covalent bonds are weak. The bonds inside molecules are strong; it is the forces between molecules that are weak (see 3.6).

Revision questions

Try these, then check your answers below each one.

Question 1. What is a covalent bond? (Choose from: The transfer of electrons from a metal to a non-metal, A shared pair of electrons between two atoms, The attraction between oppositely charged ions, A weak force between two molecules.)

Answer: A shared pair of electrons between two atoms

Working: A covalent bond is a shared pair of electrons between two atoms. Transfer of electrons makes ionic bonds.

Question 2. How many electrons are shared between the two atoms in a nitrogen molecule, N₂?

Answer: 6

Working: Each N atom shares 3 electrons, making a triple bond: 3 shared pairs = 6 electrons.

Question 3. How many lone pairs (non-bonding pairs) are on the oxygen atom in a water molecule?

Answer: 2

Working: 6 outer electrons − 2 used in O–H bonds = 4 electrons = 2 lone pairs. Oxygen then has 8 electrons in its outer shell.

Related lessons

• Previous lesson: 3.4 Giant Ionic Structures
• Next lesson: 3.6 Structure of Simple Molecules
• In this chapter: 3.1 States of Matter
• In this chapter: 3.2 Atoms and Ions
• In this chapter: 3.3 Ionic Bonding

Continue in the full course

This lesson is section 3.5 (AGCCH0305) of our free GCSE Chemistry course for AQA. In the full course you get every video for this section, a self-marking practice quiz, a printable worksheet with a mark scheme, chapter tests, and your progress saved. Go to the full GCSE Chemistry course.