Covalent bonding — AQA GCSE Chemistry

Covalent bonding describes the sharing of electron pairs between non-metal atoms seeking stable, full outer electron shells. This fundamental chemical mechanism produces either small discrete molecules such as water and oxygen or vast three-dimensional networks known as giant covalent lattices.

Mastering covalent bonding is essential across GCSE Chemistry because it explains macroscopic physical properties including states of matter, melting points, and electrical conductivity. It also lays the crucial groundwork for organic chemistry, polymers, and materials science.

• →Electrostatic Attraction: A covalent bond is the electrostatic attraction between positive atomic nuclei and a shared pair of negative electrons.
• →Dot-and-Cross Representations: Used to display how outer-shell electrons are shared, showing single, double, or triple covalent bonds.
• →Simple Molecular Structures: Molecules with strong internal covalent bonds but weak intermolecular forces between separate molecules.
• →Giant Covalent Lattices: Substances like diamond, graphite, and silicon dioxide where billions of atoms are linked continuously by strong covalent bonds.

• Covalent bonding occurs between non-metal atoms, for example hydrogen and oxygen in water.
• A covalent bond is a shared pair of electrons, with each bonded atom contributing one electron to the pair.
• The shared pair is attracted to both nuclei, which holds the atoms together in the molecule.
• A single bond is one shared pair, a double bond is two shared pairs, and a triple bond is three shared pairs.
• The number of covalent bonds an atom forms can be predicted from its group number in the periodic table.
• Covalent bonds are strong because the shared pair is attracted to both nuclei.
• A large amount of energy is needed to break covalent bonds, so they remain intact during physical changes such as melting and boiling.
• When a simple molecular substance boils, weak forces between molecules are overcome, not the strong covalent bonds within each molecule.
• Strong covalent bonds explain why molecules stay together, while weak intermolecular forces explain low melting and boiling points.
• A small molecule contains a small number of atoms joined by strong covalent bonds, such as H₂O or CO₂.
• Simple molecular substances consist of separate small molecules with weak forces between them.
• Weak intermolecular forces mean simple molecular substances have low melting and boiling points.
• Simple molecular substances do not conduct electricity because their molecules are uncharged and no free charged particles are present.
• Common small molecules can be recognised from formulas such as H₂, Cl₂, O₂, N₂, HCl, H₂O, NH₃, CH₄ and CO₂.
• A chemical formula shows the type and number of atoms in a substance.
• Small molecules contain only a few non-metal atoms, for example H₂O, CO₂, NH₃, CH₄, HCl, O₂, N₂, H₂ and Cl₂.
• Formulas such as SiO₂ represent giant covalent structures, not small molecules, even though they contain only non-metals.
• Formulas such as NaCl represent ionic compounds, not small molecules, because they contain a metal and a non-metal.
• Recognising small molecules from formulas allows prediction of low melting and boiling points and no electrical conductivity.
• Covalent bonds involve shared pairs of electrons between non-metal atoms.
• A polymer is a very large molecule built from many small monomer molecules joined by covalent bonds.
• In addition polymerisation, the double bond in each monomer opens and the monomers link into a long chain.
• The repeating unit of a polymer shows the atoms that recur along the chain, with bonds continuing on both sides.
• Polymers have high melting points because the intermolecular forces between the large polymer chains are relatively strong and require significant energy to overcome.
• A giant covalent structure is a continuous network of atoms joined by covalent bonds, not separate small molecules.
• In diamond, each carbon atom forms four single covalent bonds to other carbon atoms in a tetrahedral arrangement.
• In silicon dioxide, each silicon atom bonds to four oxygen atoms and each oxygen atom bonds to two silicon atoms.
• Melting a giant covalent substance requires breaking many strong covalent bonds, giving very high melting points.
• Diamond is very hard because its rigid network of strong covalent bonds resists deformation.
• A dot and cross diagram shows outer-shell electrons, with shared pairs in the overlap between atoms.
• A single line between two atoms represents one shared pair of electrons, and a double line represents two shared pairs.
• Ball and stick models show the arrangement of atoms in three dimensions but do not show electron pairs.
• A polymer is represented by placing the repeating unit in brackets with n outside, where n is a large number.
• The repeating unit must show the correct atoms and bonds, with bonds extending beyond the brackets to show the chain continues.
• Dot and cross diagrams must show outer-shell electrons only, with shared pairs in the overlap between atoms.
• Hydrogen, chlorine, oxygen and nitrogen are diatomic; oxygen and nitrogen contain double and triple bonds respectively.
• Water, ammonia and methane have different numbers of bonds per central atom: two, three and four respectively.
• Line representations use one line per single bond and can be used for small molecules, polymer repeating units and parts of giant covalent structures.
• Limitations include lack of shape information in dot and cross diagrams and lack of electron detail in ball and stick models.
• Molecular formula is deduced by counting the atoms of each element shown in the model or diagram.

• 💡When asked to explain bonding, name the particles involved and state that a pair of electrons is shared.
• 💡Use a dot-and-cross diagram to show one electron from each atom in the shared region, then label the bond.
• 💡Check the group number of each atom to predict how many bonds it forms before drawing a structure.
• 💡Use the phrase 'strong covalent bonds between atoms' and 'weak forces between molecules' to make the distinction clear.
• 💡When explaining a low boiling point, refer to energy needed to overcome intermolecular forces, not to break covalent bonds.
• 💡Link the strength of a covalent bond to the attraction between the shared pair and both nuclei.
• 💡Use the formula to count atoms: a small molecule contains only a few atoms, whereas a giant structure is represented by a giant lattice.
• 💡When explaining properties, always refer to weak forces between molecules rather than weak covalent bonds.
• 💡For conductivity questions, state that simple molecular substances have no free charged particles.
• 💡Check that all elements are non-metals and confirm the substance is not a known giant covalent exception (like SiO₂) before deciding it is a small molecule.
• 💡Learn a short list of common small molecules, including H₂O, CO₂, NH₃, CH₄, HCl, O₂, N₂, H₂ and Cl₂.
• 💡If a formula represents a giant covalent or ionic structure, do not describe it as a small molecule.
• 💡When asked to identify a polymer, look for a long chain or a repeating unit rather than a small fixed group of atoms.
• 💡If a diagram shows a monomer, count the carbon atoms and check whether a double bond is present before predicting the repeating unit.
• 💡Link the large size of the molecule to relatively strong intermolecular forces between the chains when explaining a high melting point, not to the covalent bonds.
• 💡When comparing diamond and silicon dioxide, state the number of bonds formed by each atom in the network.
• 💡Use the phrase giant covalent structure rather than molecule when describing diamond or silicon dioxide.
• 💡Explain high melting point by referring to the many strong covalent bonds that must be broken, not to weak intermolecular forces.
• 💡When drawing a polymer, check that the bonds at each end of the repeating unit pass through the brackets.
• 💡Use a line for each single bond and two parallel lines for a double bond, and keep the number of lines consistent with the atoms shown.
• 💡If asked what n represents, state that it is a large number of repeating units, not a specific value.
• 💡Check the group number of each atom to work out how many covalent bonds it forms before drawing.
• 💡When deducing a molecular formula, count atoms systematically and use subscript numbers for more than one atom.
• 💡For limitations, name one specific feature each model cannot show rather than saying it is simply inaccurate.
• 💡Always specify whether you are referring to 'intermolecular forces' or 'covalent bonds' in 4-mark and 6-mark bonding answers.
• 💡When drawing dot-and-cross diagrams, ensure non-bonding outer-shell electrons (lone pairs) are included unless the question explicitly instructs you to show bonding pairs only.
• 💡Memorise the exact definitions of single, double, and triple bonds in terms of the number of shared pairs of electrons (1, 2, and 3 pairs respectively).

• Thinking that atoms transfer electrons in covalent bonding; correct this by stating that electrons are shared, whereas transfer produces ions.
• Drawing only one electron in a shared pair; correct this by showing both electrons in the overlap region between the atoms.
• Assuming all covalent bonds are single; correct this by checking the formula, for example O₂ contains a double bond made of two shared pairs.
• Saying that covalent bonds break when a simple molecular substance boils; correct this by stating that intermolecular forces break instead.
• Confusing strong covalent bonds with weak intermolecular forces; correct this by naming each type of attraction separately.
• Claiming that all attractions in a simple molecular substance are weak; correct this by stating that the bonds within each molecule are strong.
• Thinking that small molecules contain weak covalent bonds; correct this by stating that the covalent bonds are strong but the forces between molecules are weak.
• Believing that simple molecular substances conduct electricity when molten; correct this by stating that they contain no free ions or electrons.
• Confusing a small molecule with a giant covalent structure; correct this by checking whether the formula represents a few atoms only.
• Treating every formula containing non-metals as a small molecule; correct this by memorising exceptions like SiO₂ which is a giant covalent structure.
• Assuming that a formula with two elements is always a small molecule; correct this by identifying ionic compounds such as NaCl.
• Counting atoms incorrectly in formulas with subscripts; correct this by multiplying the subscript by any number outside a bracket, as in Ca(OH)₂.
• Thinking a polymer is a mixture of separate small molecules; correct this by describing one very large molecule made of many repeating units joined by covalent bonds.
• Drawing the polymer repeating unit with a double bond still present; correct this by showing single bonds only, because the double bond opens during polymerisation.
• Attributing the high melting point of polymers to breaking covalent bonds; correct this by stating that melting overcomes the relatively strong intermolecular forces between the chains, not the covalent bonds.
• Describing diamond as a simple molecular substance; correct this by stating that diamond is a giant covalent network with no separate molecules.
• Saying silicon dioxide contains Si₂O₄ molecules; correct this by describing a continuous network in which each silicon atom is bonded to four oxygen atoms.
• Claiming giant covalent substances melt easily because covalent bonds are weak; correct this by explaining that many strong covalent bonds must be broken, so melting points are very high.
• Drawing a polymer repeating unit without brackets or without n; correct this by enclosing the repeating unit in brackets and writing n outside.
• Showing a double bond inside a polymer repeating unit; correct this by drawing single bonds only, because the double bond opens during polymerisation.
• Treating n as a small fixed number such as 2 or 3; correct this by stating that n is a large number of repeating units.
• Showing all electrons rather than only outer-shell electrons in a dot and cross diagram; correct this by drawing all outer-shell electrons, including both shared bonding pairs and non-bonding lone pairs.
• Drawing oxygen with a single bond; correct this by showing a double bond, because oxygen atoms share two pairs of electrons.
• Giving the molecular formula as the empirical formula or missing subscripts; correct this by counting every atom shown, for example C₂H₆ rather than CH₃.
• Believing that covalent bonds break when simple molecular substances boil or melt, rather than weak intermolecular forces.
• Assuming all covalent substances have low melting points, overlooking giant covalent structures like quartz and diamond.
• Thinking that covalent substances conduct electricity when dissolved in water, failing to recognise that neutral molecules do not dissociate into mobile ions (with the exception of acids).

1. 1Day 1-2: Master dot-and-cross diagrams for common diatomic molecules (H2, Cl2, O2, N2, HCl) and simple compounds (H2O, NH3, CH4).
2. 2Day 3-4: Practice writing standard comparative explanations for the melting and boiling points of simple molecular substances versus giant covalent structures.
3. 3Day 5: Revise molecular models (dot-and-cross, 2D structural formulas, 3D ball-and-stick) and evaluate their strengths and limitations.
4. 4Day 6-7: Complete AQA past paper 6-mark questions on bonding, structures, and physical properties under timed conditions.

• 📋Dot-and-cross diagram completion: Drawing shared pairs and lone pairs for molecules like ammonia or water.
• 📋Explain physical properties: Explaining why a named simple molecular substance has a low boiling point or does not conduct electricity.
• 📋Evaluate molecular models: Stating one advantage and one limitation of 2D displayed formulas or 3D ball-and-stick representations.