In simple terms
A friendly intro before the formal notes — no formulas yet.
Sharing, Shaping, Sticking
Non-metal atoms bond by SHARING electron pairs — that is the covalent bond. The number of pairs shared sets how strong and short the bond is; the arrangement of pairs around an atom sets the molecule's SHAPE; and the way whole molecules attract each other sets how they behave in bulk, such as their boiling point.
Think of two people gripping the same rope. One shared grip (a single bond) holds them together; gripping with both hands each (a triple bond) pulls them much closer and is far harder to prise apart. Now imagine several ropes tied to one central person — they naturally push as far apart as possible, and that spacing is the molecule's shape. Finally, whole groups of rope-linked people can be drawn towards other groups by weaker tugs between them; those weak tugs are the intermolecular forces that decide whether the substance is a gas, liquid or solid.
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See the bond: a covalent bond is a shared pair of electrons attracted to both nuclei. Count how many pairs are shared to get the bond order.
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Draw the molecule: use a Lewis structure to place every bonding pair and lone pair, so you know what surrounds the central atom.
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Find the shape: count the electron domains (bonds and lone pairs) around the central atom; VSEPR says they spread out as far as possible, and lone pairs squeeze the bond angles.
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Predict the bulk behaviour: decide if the molecule is polar, then rank the intermolecular forces (dispersion < dipole-dipole < hydrogen bonding) to explain boiling point and solubility.
Explore the concept
Use the live diagram, PhET or GeoGebra sim, and synced steps — play it, drag controls, or tap a step.
Step 1
See the bond: a covalent bond is a shared pair of electrons attracted to both nuclei. Count how many pairs are shared to get the bond order.
Full topic notes
Formal explanation with the rigour you need for the exam.
What a covalent bond is
A covalent bond is the electrostatic attraction between a shared pair of electrons and the positively charged nuclei of the two atoms sharing them. Atoms share so that each reaches a stable, full outer shell — an octet of eight valence electrons for most atoms (the octet rule), or a duet of two for hydrogen, matching the nearest noble gas. Because the shared electrons sit between the two nuclei, both nuclei are attracted to the same region of negative charge, and that mutual attraction holds the atoms together.
A covalent bond = a shared pair of electrons attracted to two nuclei; it forms between non-metals.
Atoms share to reach a full outer shell (octet, or a duet for H).
A bonding pair is shared between two atoms; a lone pair belongs to one atom and is not shared.
Boron (e.g. BF₃, 6 electrons) and beryllium (e.g. BeCl₂, 4 electrons) are common incomplete-octet exceptions at SL.
Bond order: single, double and triple bonds
Atoms can share one, two or three pairs of electrons, giving a single, double or triple bond respectively. The number of shared pairs is the bond order, and it directly controls two measurable properties. More shared pairs mean more negative charge concentrated between the nuclei, so the nuclei are pulled closer (shorter bond) and held more tightly (stronger bond, higher bond enthalpy).
Single bond (order 1): longest and weakest, e.g. C–C.
Double bond (order 2): intermediate length and strength, e.g. C=C.
Triple bond (order 3): shortest and strongest, e.g. C≡C and the N≡N of nitrogen gas, whose great strength makes N₂ very unreactive.
Drawing Lewis (electron-dot) structures
A Lewis structure is a 2D map of a molecule that shows every bonding pair (as a line) and every lone pair (as a pair of dots). It is the tool you use before deducing shape and polarity, so getting it right is the foundation of the whole topic. Follow the same four steps every time.
Step 1 — Count valence electrons. Sum the valence electrons of all atoms. For a polyatomic ion, ADD one electron per negative charge and SUBTRACT one per positive charge.
Step 2 — Arrange the atoms. Put the least electronegative atom in the centre. Hydrogen and halogens are almost always terminal (on the outside).
Step 3 — Single bonds, then lone pairs. Join the central atom to the outer atoms with single bonds, then place remaining electrons as lone pairs on the outer atoms first until they have an octet.
Step 4 — Make multiple bonds if needed. If the central atom lacks an octet, convert outer lone pairs into double or triple bonds until it is satisfied.
Bond polarity and electronegativity
When the two atoms in a bond have different electronegativities, the more electronegative atom pulls the shared pair closer to itself, gaining a partial negative charge (δ−) while the other becomes partial positive (δ+). This separation of charge is a polar bond, and it is described by a bond dipole pointing towards the δ− atom. The larger the electronegativity difference, the more polar the bond; if the difference is zero (two identical atoms, as in Cl₂ or O₂) the bond is non-polar. A very large difference tips the bonding towards ionic — bonding is a continuum, not a set of separate boxes.
Electronegativity increases across a period and decreases down a group; fluorine is the most electronegative element.
Bigger electronegativity difference → more polar bond → larger δ+/δ− and a bigger bond dipole.
Equal electronegativity (identical atoms) → non-polar bond.
A polar bond is a necessary condition for a polar molecule, but as we will see it is not sufficient — the shape matters too.
VSEPR: predicting molecular shape and bond angle
Valence Shell Electron Pair Repulsion (VSEPR) theory says that the electron domains around a central atom — each single, double or triple bond counts as ONE domain, and each lone pair counts as one domain — arrange themselves as far apart as possible, because negatively charged clouds repel. Count the domains, place them for maximum separation, then describe the shape using only the ATOM positions. Lone pairs occupy space and set the geometry, but they are invisible in the named shape.
2 domains, 0 lone pairs → linear, 180° (e.g. CO₂, BeCl₂).
3 domains, 0 lone pairs → trigonal planar, 120° (e.g. BF₃, CO₃²⁻).
4 domains, 0 lone pairs → tetrahedral, 109.5° (e.g. CH₄, CCl₄).
4 domains, 1 lone pair → trigonal pyramidal, ~107° (e.g. NH₃).
4 domains, 2 lone pairs → bent / V-shaped, ~104.5° (e.g. H₂O).
Each lone pair repels more strongly than a bonding pair, so it squeezes the remaining bond angles smaller by roughly 2–2.5°.
From shape to molecular polarity
A molecule is polar if it has an overall dipole — a net separation of charge across the whole molecule. To decide, check two things: are the bonds polar, and does the shape let the bond dipoles cancel? If the shape is symmetric (the polar bonds point in balanced, opposing directions), the dipoles cancel and the molecule is non-polar overall, even though its bonds are polar. If the shape is asymmetric, usually because of a lone pair or unlike terminal atoms, the dipoles do not cancel and the molecule is polar.
Polar bonds + symmetric shape → non-polar molecule: CO₂ (linear), BF₃ (trigonal planar), CH₄ and CCl₄ (tetrahedral) — dipoles cancel.
Polar bonds + asymmetric shape → polar molecule: H₂O (bent) and NH₃ (trigonal pyramidal) — dipoles reinforce because of the lone pair(s).
A polar molecule has a permanent dipole and so can attract other molecules through dipole-dipole forces.
This is the single most common trap in the topic: never conclude a molecule is polar just because it contains polar bonds — always check the shape.
Intermolecular forces: boiling points and solubility
The covalent bonds we have discussed are the strong forces INSIDE a molecule. What decides whether a molecular substance is a gas, liquid or solid — and its boiling point and solubility — are the much weaker forces BETWEEN molecules. When a molecular substance boils you overcome these intermolecular forces, not the covalent bonds, so ranking them correctly is the key skill. There are three, in increasing strength.
London (dispersion) forces: present between ALL molecules. They arise from temporary, instantaneous dipoles as electrons move, and get stronger as the number of electrons (molar mass) increases. They are the ONLY force between non-polar molecules.
Dipole-dipole forces: additional attractions between the permanent dipoles of polar molecules (δ+ to δ−). Stronger than dispersion for molecules of similar size, so a polar molecule boils higher than a non-polar one of similar mass.
Hydrogen bonding: the strongest of the three. It requires H bonded directly to N, O or F, attracted to a lone pair on an N, O or F of a neighbouring molecule. It explains the anomalously high boiling points of H₂O, NH₃ and HF.
Solubility: 'like dissolves like'. Polar and hydrogen-bonding substances dissolve in water (e.g. ethanol, ammonia); non-polar substances do not, but dissolve in non-polar solvents.
Giant covalent (network) structures
Not every covalent substance is made of small molecules. In giant covalent (network) solids the covalent bonds extend continuously through the whole structure, so there are no separate molecules and no weak forces to overcome on melting — you must break strong covalent bonds throughout. This gives very high melting points and distinctive properties that come straight from how the atoms are connected.
Diamond: each carbon bonds covalently to four others in a rigid 3D tetrahedral network. Extremely hard, very high melting point, and does NOT conduct electricity (all four outer electrons are localised in bonds).
Graphite: each carbon bonds to only three others in flat hexagonal layers; the layers are held together by weak London forces so they slide (a lubricant, soft). The fourth electron per carbon is delocalised between layers, so graphite CONDUCTS electricity.
Graphene: a single one-atom-thick layer of graphite — an extremely strong, flexible sheet that conducts electricity for the same delocalised-electron reason.
Silicon dioxide (SiO₂, quartz): each Si bonds to four O and each O bridges two Si in a giant 3D network. Very high melting point, hard, and an electrical insulator (no delocalised electrons).
Common mistakes examiners penalise
Calling a molecule polar just because its bonds are polar — always check the shape; symmetric molecules like CO₂, CCl₄ and BF₃ have polar bonds but cancel to a non-polar molecule.
Forgetting lone pairs when counting VSEPR domains — a lone pair is an electron domain and changes the shape and angle; NH₃ is pyramidal (not trigonal planar) and H₂O is bent (not linear) because of their lone pairs.
Quoting 109.5° for NH₃ or H₂O — lone-pair repulsion shrinks the angle to ~107° and ~104.5° respectively; you must state the lone-pair reasoning, not just the number.
Claiming hydrogen bonding without H bonded to N, O or F — molecules like CH₃F and HCl are polar and use dipole-dipole forces, but they cannot hydrogen-bond because the H is not on N, O or F.
Saying covalent bonds break when a molecular substance boils — boiling overcomes the weak intermolecular forces; the covalent bonds inside each molecule stay intact.
Ignoring dispersion forces because a molecule is non-polar — dispersion is present in everything and is the reason non-polar substances like the halogens have rising boiling points down the group.
Confusing diamond and graphite conductivity — graphite conducts (delocalised electrons, 3 bonds per C); diamond does not (4 bonds per C, all electrons localised).
Model answer — marked the way our engine marks it
This is the showcase for a structured Paper 2 question. The marks are analytic, not a markband: each mark is tied to a specific point of chemistry — a method/reasoning mark (M) for correct reasoning, an answer mark (A) for a correct shape, angle or conclusion — and error-carried-forward (ECF) means a wrong deduction early on can still earn the marks that follow if you use it correctly. Any chemically equivalent wording is accepted. Study how each of the four marks below is earned by one distinct point.
Where this leads
The covalent model underpins much of what follows. Organic chemistry is built on the shapes, polarities and reactivity of covalently bonded carbon; the properties of water — its solvent power, its high boiling point, its role in life — trace straight back to hydrogen bonding and molecular polarity; and energetics uses bond enthalpies, which are just the bond strengths introduced here. Master the chain — Lewis structure, then VSEPR shape and angle, then polarity, then intermolecular forces — and you can predict the physical behaviour of almost any molecular substance in the course.
Worked examples
See the formulas applied — reveal one step at a time, like the exam.
Draw the Lewis structure of carbon dioxide, CO₂, and state the carbon–oxygen bond order. [3]
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Step 1 — count valence electrons. Carbon (Group 14) has 4; each oxygen (Group 16) has 6. Total electrons. [M1: correct electron count]
Deduce the molecular shape and the H–O–H bond angle in a water molecule, H₂O, and explain the value of the angle. [4]
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Step 1 — electron domains on the central atom. Oxygen (Group 16) has 6 valence electrons. Two form single bonds to the two hydrogens (2 bonding pairs); the other four are two lone pairs. Total domains . [M1: correct count of bonding pairs and lone pairs]
The three substances ethane (C₂H₆, Mᵣ ≈ 30), fluoromethane (CH₃F, Mᵣ ≈ 34) and methanol (CH₃OH, Mᵣ ≈ 32) have similar molar masses but very different boiling points: −89 °C, −78 °C and +65 °C. Assign each boiling point and explain the order in terms of intermolecular forces. [4]
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Step 1 — identify the strongest force in each.
- Ethane, C₂H₆: non-polar, so only London dispersion forces.
- Fluoromethane, CH₃F: polar (asymmetric, polar C–F bond) but no H on N/O/F, so dipole-dipole plus dispersion.
- Methanol, CH₃OH: has O–H, so hydrogen bonding plus dipole-dipole and dispersion. [M1: correct force identified for each]
Deduce the shape and bond angle of ammonia, NH₃, and explain why they differ from those of methane, CH₄. [4]
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Model answer — full reasoning.
How it all connects
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Tap a linked idea to see how it connects back to the main topic — that connection is what examiners reward.
Glossary
Try to recall each definition before you reveal it.
Quick check
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Revision flashcards
Flip the card. Test yourself before the exam.
Covalent bond
The electrostatic attraction between a shared pair of electrons and the positively charged nuclei of the two bonded atoms. It forms between non-metals.
Key takeaways
Review these before you close the topic — retrieval beats re-reading.
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A covalent bond = a shared pair of electrons attracted to two nuclei; it forms between non-metals.
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Atoms share to reach a full outer shell (octet, or a duet for H).
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A bonding pair is shared between two atoms; a lone pair belongs to one atom and is not shared.
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Boron (e.g. BF₃, 6 electrons) and beryllium (e.g. BeCl₂, 4 electrons) are common incomplete-octet exceptions at SL.
Practice — then mark it
The whole point: a real Cambridge question, marked mark-by-mark.
Get a Paper 2 structured question marked: deduce a shape and bond angle, or rank boiling points by intermolecular force, with full reasoning
Get a Paper 2 structured question marked: deduce a shape and bond angle, or rank boiling points by intermolecular force, with full reasoning
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Frequently asked
Checkpoint
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Reading it isn’t knowing it — prove it.
Before you move on: do Get a Paper 2 structured question marked: deduce a shape and bond angle, or rank boiling points by intermolecular force, with full reasoning on paper, snap a photo, and get examiner-style feedback on exactly where you win and lose marks.