Revision notes for AQA GCSE Chemistry Covalent bonding. Open the guide for explanations and worked examples. Written against the AQA GCSE Chemistry (8462) specification, so the content matches what's examinable rather than general Chemistry background.
Covalent bonding
What you'll learn
- What a covalent bond is and why it forms between non-metal atoms.
- How to draw dot-and-cross diagrams for common GCSE molecules.
- How to use lines to represent single, double and triple covalent bonds.
- How to recognise small molecules, polymers and giant covalent structures from formulas and diagrams.
The key prerequisite: outer electrons
Electrons are arranged in shells around an atom. The outer shell is the furthest shell from the nucleus that contains electrons. In bonding, the outer-shell electrons are the important ones.
For the GCSE molecules in this topic, atoms often bond so that:
- hydrogen has 2 electrons in its outer shell
- carbon, nitrogen, oxygen and chlorine have 8 electrons in their outer shell
This is a stable arrangement, like the electron arrangement of a noble gas.
Tip
Outer electrons from the periodic table
For many GCSE examples, the group number tells you the number of outer-shell electrons: carbon has 4, nitrogen has 5, oxygen has 6, chlorine has 7, and hydrogen has 1.
What is a covalent bond?
A non-metal is an element that does not show typical metallic properties; non-metals are found mainly on the right-hand side of the periodic table, with hydrogen as a special case on the left.
When non-metal atoms bond with other non-metal atoms, they usually share electrons rather than transfer them.
Definition
Covalent bond
A covalent bond is a strong bond formed when atoms share a pair of electrons. The shared pair is attracted to the positive centres of both atoms.
One shared pair of electrons forms one single covalent bond. Two shared pairs form a double covalent bond, and three shared pairs form a triple covalent bond.
Key Idea
Shared means counted by both atoms
In covalent bonding, an electron pair is shared. Each bonded atom counts the shared electrons as part of its own outer shell.
Common Mistake
Thinking electrons are transferred
Covalent bonding is not one atom giving electrons away. That is ionic bonding. In covalent bonding, atoms share pairs of electrons.
Molecules
A molecule is a particle made from two or more atoms joined by covalent bonds. Many covalently bonded substances are made of small molecules, meaning each molecule contains a small, fixed number of atoms.
Common GCSE examples include:
- hydrogen, H₂
- chlorine, Cl₂
- oxygen, O₂
- nitrogen, N₂
- hydrogen chloride, HCl
- water, H₂O
- ammonia, NH₃
- methane, CH₄
The formula tells you how many atoms are in one molecule. For example, H₂O means each water molecule contains 2 hydrogen atoms and 1 oxygen atom.
Dot-and-cross diagrams
A dot-and-cross diagram shows the outer-shell electrons in a covalent molecule. Dots and crosses are used to show which atom each electron came from.
Only the outer-shell electrons are shown. Inner-shell electrons are usually left out because they are not involved in bonding.
A lone pair is a pair of outer-shell electrons that is not shared in a covalent bond.
The diagram below shows common dot-and-cross diagrams for hydrogen, hydrogen chloride, water, ammonia and methane.
Common Mistake
Drawing all the electrons
For dot-and-cross diagrams at GCSE, draw the outer-shell electrons only. Do not draw every electron in the atom unless the question specifically asks for full electronic structures.
A method for drawing dot-and-cross diagrams
Use this slow, reliable method:
- Find the number of outer-shell electrons for each atom.
- Decide how many electrons each atom needs for a full outer shell.
- Put atoms close enough that their outer shells overlap.
- Put one electron from each atom into each overlap to make a shared pair.
- Add any remaining outer electrons as lone pairs.
- Check that hydrogen has 2 electrons and the other common atoms have 8.
Example
Drawing a dot-and-cross diagram for water
- Oxygen is in Group 6, so it has 6 outer-shell electrons. Each hydrogen atom has 1 outer-shell electron, so there are 8 outer-shell electrons to place in total.
- Put oxygen in the centre because it needs 2 more electrons to complete its outer shell, and each hydrogen needs 1 more electron.
- Make two shared pairs: one shared pair between oxygen and the first hydrogen, and one shared pair between oxygen and the second hydrogen.
- Count the shared electrons for both atoms. Each hydrogen now has 2 electrons, and oxygen now has 8 electrons.
- Place oxygen’s remaining 4 electrons as two lone pairs.
Single, double and triple bonds
A structural formula shows how atoms are connected, using lines for covalent bonds. One line represents one shared pair of electrons.
So:
- H–H has one single covalent bond
- Cl–Cl has one single covalent bond
- H–Cl has one single covalent bond
- O=O has one double covalent bond
- N≡N has one triple covalent bond
The diagram below links dot-and-cross diagrams to structural formulas for chlorine, oxygen and nitrogen.
Tip
Bond-count shortcut
For common GCSE molecules: hydrogen and chlorine usually form 1 bond, oxygen usually forms 2 bonds, nitrogen usually forms 3 bonds, and carbon usually forms 4 bonds.
Representing common molecules with lines
Line diagrams are quicker than dot-and-cross diagrams, but they show less detail.
Common examples:
- hydrogen: H–H
- chlorine: Cl–Cl
- hydrogen chloride: H–Cl
- water: H–O–H
- ammonia: nitrogen bonded to 3 hydrogen atoms
- methane: carbon bonded to 4 hydrogen atoms
- oxygen: O=O
- nitrogen: N≡N
Common Mistake
Forgetting multiple bonds
O₂ is not O–O in the usual GCSE covalent diagram; it is O=O. N₂ is not N–N; it is N≡N.
Deducing a molecular formula from a model
A molecular formula tells you the actual number of each type of atom in one molecule. To find it from a model, count atoms, not bonds.
Example
Deducing a formula from a molecular model
A model shows one carbon atom bonded to one oxygen atom by a double bond, and bonded to two hydrogen atoms by single bonds.
- Count the atoms shown in the model: there is 1 carbon atom, 2 hydrogen atoms and 1 oxygen atom.
- Write the symbols with the numbers of atoms: C, H₂ and O.
- Combine them into the molecular formula: CH₂O. The double bond changes how atoms are connected, but it does not change how many atoms are present.
Different covalent structures
Covalently bonded substances are not all small molecules. They can be grouped into three main types.
A small molecule has a small, fixed number of atoms, such as H₂O, NH₃ or CH₄.
A polymer is a very large molecule made from many repeating units joined together. A repeating unit is the small section that appears again and again in the chain. For example, poly(ethene) can be represented as [-CH₂-CH₂-]ₙ, where n is a large number.
A giant covalent structure is a huge network of atoms joined by covalent bonds throughout the structure. Diamond and silicon dioxide, SiO₂, are examples. There are no separate small molecules in a giant covalent structure.
The diagram below compares small molecules, polymers and giant covalent structures.
Example
Classifying covalent structures
A diagram shows three substances: separate CO₂ particles, a long chain with brackets and a subscript n, and a continuous network of carbon atoms.
- The separate CO₂ particles are small molecules because each particle has a fixed number of atoms.
- The long chain with brackets and n is a polymer because it shows a repeating unit repeated many times.
- The continuous carbon network is a giant covalent structure because the covalent bonds extend throughout the structure rather than stopping at separate molecules.
Common Mistake
Formulas do not always show molecules
SiO₂ is a formula for silicon dioxide, but silicon dioxide is a giant covalent structure, not separate SiO₂ molecules. In giant structures, the formula shows the ratio of atoms, not a single molecule.
Limitations of different diagrams
Every model leaves something out. In exams, you may be asked to describe limitations.
Dot-and-cross diagrams
Dot-and-cross diagrams are useful because they show which electrons are shared and which are lone pairs.
Limitations:
- They do not show the real 3D shape of the molecule.
- They do not show the real sizes of atoms.
- Dots and crosses are just symbols; real electrons are identical.
Structural formulas and 2D diagrams
A 2D diagram is drawn flat on the page. It is useful for showing which atoms are bonded together.
Limitations:
- It may make a 3D molecule look flat.
- It often leaves out lone pairs.
- It does not show the real bond angles.
Ball-and-stick models
A ball-and-stick model represents atoms as balls and bonds as sticks.
Limitations:
- Atoms are not really joined by sticks.
- The sizes of atoms and bond lengths may not be accurate.
- It can exaggerate the spaces between atoms.
3D and space-filling models
A 3D model shows depth and shape more realistically. A space-filling model shows atoms as overlapping spheres, closer to the space atoms actually occupy.
Limitations:
- Bonds can be harder to see.
- Atoms at the back may be hidden.
- It still simplifies the real arrangement of electrons.
Exam technique
In the exam
- For dot-and-cross diagrams, count outer-shell electrons first, then check that hydrogen has 2 electrons and common atoms like carbon, nitrogen, oxygen and chlorine have 8.
- For line diagrams, remember that one line is one shared pair: O₂ needs a double bond and N₂ needs a triple bond.
- When classifying structures, look for separate fixed molecules, a bracketed repeating unit with n, or a continuous network of covalent bonds.
Self review
Check yourself
- How many shared pairs of electrons are there in O₂ and in N₂?
- In a dot-and-cross diagram of water, how many lone pairs are on the oxygen atom?
- What feature in a diagram tells you that a covalent substance is a polymer rather than a small molecule?
Recap questions
Test yourself with 5 quick questions on this guide. Answer them all correctly to complete it.
Chemical bonds, ionic, covalent and metallic
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