What you'll learn
Covalent bonding and simple molecular substances is the second bonding type in AQA GCSE Chemistry, and it contains the single distinction that generates more marks than any other in the bonding unit: the difference between the strong covalent bonds inside a molecule and the weak intermolecular forces between molecules. Almost every question about melting point, boiling point or state of matter in this topic turns on that distinction, and almost every wrong answer confuses the two. Covalent bonding occurs between non-metals, which share pairs of electrons rather than transferring them. By the end of this guide you should be able to explain how covalent bonds form, work out how many bonds an atom will make, represent molecules in several ways, explain why simple molecular substances have low melting points, account for the trend with molecular size, explain why they do not conduct electricity, and distinguish simple molecular from giant covalent substances.
Key terms and definitions
Covalent bond — a shared pair of electrons between two non-metal atoms
Molecule — a discrete group of atoms held together by covalent bonds
Simple molecular substance — a substance made of small separate molecules
Intermolecular force — the weak force of attraction acting between separate molecules
Intramolecular bond — a bond within a molecule, which in this topic means the covalent bond
Single bond — one shared pair of electrons
Double bond — two shared pairs of electrons between the same two atoms
Lone pair — a pair of outer electrons not involved in bonding
Displayed formula — a representation showing every atom and every bond as a line
Diatomic molecule — a molecule containing two atoms, such as hydrogen or chlorine
Volatile — easily evaporated, having a low boiling point
Core concepts
How covalent bonds form
Covalent bonding occurs between non-metal atoms. Rather than transferring electrons, as happens between a metal and a non-metal, the atoms share pairs of electrons so that each achieves the electronic structure of a noble gas.
In a hydrogen molecule, each atom contributes one electron to a shared pair, and both atoms effectively have the two electrons of helium.
In a chlorine molecule, each atom contributes one electron, and both effectively have eight outer electrons, matching argon.
The covalent bond itself is strong. It results from the electrostatic attraction between the shared pair of negative electrons and the positive nuclei of both atoms, and a large amount of energy is needed to break it. Remembering that covalent bonds are strong is essential, because the low melting points of molecular substances are so often wrongly attributed to weak covalent bonds.
Predicting the number of bonds
The number of covalent bonds an atom forms equals the number of electrons it needs to complete its outer shell, which is eight minus its group number.
Group 4 elements such as carbon form four bonds. Group 5 elements such as nitrogen form three. Group 6 elements such as oxygen form two. Group 7 elements such as chlorine form one. Hydrogen is the exception, forming one bond because it needs only two electrons in total.
This rule predicts molecular formulae reliably. Carbon needing four bonds and hydrogen one gives methane with four hydrogens; nitrogen needing three gives ammonia with three; oxygen needing two gives water with two.
Double bonds
Where two atoms share two pairs of electrons, a double bond forms. Oxygen is the standard example: each atom needs two more electrons, so the two atoms share two pairs.
Carbon dioxide contains two double bonds, one between the carbon and each oxygen, which satisfies carbon's requirement for four bonds and each oxygen's for two.
A double bond is stronger and shorter than a single bond, and it is also the reactive site in an alkene, which links this topic to organic chemistry.
Representing molecules
Several representations are used and each has strengths and limitations that are themselves examinable.
A dot and cross diagram shows the outer electrons of each atom using different symbols, making clear which atom each electron came from. It shows electron arrangement well but gives no sense of shape or relative size.
A displayed formula shows every atom with each covalent bond drawn as a straight line. It shows which atoms are joined to which, but represents electrons only implicitly and does not show the true three-dimensional shape.
A molecular formula gives only the number of each type of atom, which is compact but shows nothing about structure.
A ball and stick model shows the three-dimensional arrangement and bond angles, but exaggerates the space between atoms.
A space-filling model shows relative size and shape realistically but obscures the bonds.
When a question asks you to evaluate a representation, the answer generally concerns what it shows against what it omits.
Why simple molecular substances have low melting points
This is the central explanation of the topic and it must be stated precisely.
Substances made of small molecules — hydrogen, oxygen, nitrogen, chlorine, water, ammonia, methane, carbon dioxide — have low melting and boiling points, and many are gases or liquids at room temperature.
When such a substance melts or boils, only the weak intermolecular forces between the molecules are overcome. The strong covalent bonds within each molecule are not broken at all; the molecules simply move apart from one another, remaining intact.
Because the intermolecular forces are weak, very little energy is needed, so the melting and boiling points are low.
The sentence that earns full marks always contains three elements: the forces overcome are intermolecular, those forces are weak, and the covalent bonds remain unbroken.
The trend with molecular size
As molecules get larger, the melting and boiling points increase.
The reason is that intermolecular forces become stronger between larger molecules, so more energy is required to separate them.
This trend explains a great deal elsewhere in the course. It explains why the hydrocarbon fractions from crude oil separate by boiling point, why larger hydrocarbons are liquids and solids while smaller ones are gases, and why polymers are solids at room temperature despite being covalently bonded — polymer molecules are so large that the intermolecular forces between them are relatively strong.
Why they do not conduct electricity
Simple molecular substances do not conduct electricity in any state.
The molecules have no overall electrical charge, and all the outer electrons are either in bonding pairs or in lone pairs, so there are no delocalised electrons and no free ions. With no charged particles free to move, no current can flow.
This holds even when the substance is melted, which distinguishes simple molecular substances from ionic compounds.
Simple molecular against giant covalent
Both involve covalent bonds, and confusing them is common, so the distinction is worth stating explicitly.
A simple molecular substance consists of small discrete molecules with weak forces between them, giving low melting points.
A giant covalent structure has covalent bonds extending throughout the entire structure, with no separate molecules and therefore no intermolecular forces, giving very high melting points. Diamond, graphite and silicon dioxide are the examples.
The diagnostic is melting point. A covalently bonded substance melting below about 100 degrees Celsius is simple molecular; one melting above several hundred degrees is giant covalent.
Note that fullerenes such as C60, despite being made of carbon and containing many atoms, are simple molecular, because they exist as separate molecules.
Worked examples
Example 1: Predicting a molecular formula (3 marks)
Nitrogen is in Group 5 and hydrogen forms one bond. Deduce the formula of the compound they form and explain your reasoning.
A nitrogen atom has five outer electrons and therefore needs three more to complete its outer shell, so it forms three covalent bonds. A hydrogen atom needs one more electron and forms one bond.
Three hydrogen atoms are therefore needed to satisfy nitrogen's three bonds, giving ammonia with one nitrogen and three hydrogen atoms. Each shared pair gives nitrogen a full outer shell of eight and each hydrogen a full shell of two.
Example 2: Explaining a low boiling point (4 marks)
Chlorine boils at minus 34 degrees Celsius even though the covalent bond within a chlorine molecule is strong. Explain why.
Chlorine consists of small separate molecules, each containing two atoms joined by a strong covalent bond. Boiling requires only the forces between the molecules to be overcome, not the bonds within them, so the covalent bonds remain intact and the molecules simply move apart.
The intermolecular forces between chlorine molecules are weak, so only a small amount of energy is needed to overcome them. The boiling point is therefore very low, and chlorine is a gas at room temperature.
Example 3: Classifying from data (4 marks)
Substance X melts at minus 95 degrees Celsius and does not conduct electricity in any state. Substance Y melts at 1,610 degrees Celsius and does not conduct electricity in any state. Deduce the structure of each.
Substance X has a very low melting point, which shows that only weak forces need to be overcome. This indicates small separate molecules with weak intermolecular forces between them, so X is a simple molecular substance. Its failure to conduct is consistent, since molecules carry no charge and there are no delocalised electrons.
Substance Y has a very high melting point, which shows that many strong bonds must be broken, indicating a giant structure. Because it does not conduct even when molten, it cannot be ionic, and because it does not conduct as a solid it cannot be metallic. Y therefore has a giant covalent structure.
Common mistakes and how to avoid them
The most heavily penalised error in the entire bonding unit is stating that covalent bonds are broken when a simple molecular substance melts or boils. The intermolecular forces are overcome; the covalent bonds survive.
Closely related is describing intermolecular forces as weak covalent bonds. They are not bonds at all, and the phrasing suggests the misconception the examiner is testing for.
Students often say that a substance does not conduct because it has no ions, omitting that it also has no delocalised electrons. Both are usually needed.
Another common slip is assuming that a large number of atoms means a giant structure. A C60 molecule contains sixty atoms and is still simple molecular, because the molecules are separate.
Finally, many answers explain the size trend by saying larger molecules have stronger covalent bonds. The covalent bonds are irrelevant to boiling; it is the intermolecular forces that increase with size.
Exam technique for "Covalent bonding and simple molecular substances"
Check the elements first. Two or more non-metals means covalent, which settles the bonding type before any reasoning begins.
When explaining melting or boiling points, write the word intermolecular explicitly and state that the covalent bonds are not broken. Examiners look for both.
In dot and cross diagrams, draw only the outer shell electrons and use different symbols for the two atoms. Count the electrons around each atom before finishing to confirm every atom has a full outer shell.
For classification questions, use melting point to separate simple molecular from giant, then conductivity in the solid and molten states to separate giant covalent from ionic and metallic.
Quick revision summary
Covalent bonding occurs between non-metals, which share pairs of electrons so that each achieves a noble gas structure, and the covalent bond itself is strong. The number of bonds an atom forms is eight minus its group number, with hydrogen forming one, so carbon forms four, nitrogen three, oxygen two and the halogens one; sharing two pairs gives a double bond, as in oxygen and carbon dioxide. Simple molecular substances have low melting and boiling points because only the weak intermolecular forces between molecules are overcome, while the strong covalent bonds within each molecule remain intact. Melting and boiling points rise as molecules get larger, because intermolecular forces increase with size, which explains fractional distillation and why polymers are solids. Simple molecular substances never conduct, because the molecules carry no charge and there are no delocalised electrons or free ions. Giant covalent structures differ in having covalent bonds throughout with no separate molecules, giving very high melting points — but fullerenes remain simple molecular despite their size.