From C2 Bonding & Structure you already know ionic, covalent and metallic bonding, dot-and-cross diagrams, and the giant/simple structure split. A-Level keeps all of it and upgrades each part: bonds gain a third kind (dative), “intermolecular forces” splits into three named forces you must tell apart, shapes get exact names and angles you can derive, and electronegativity turns “polar” from a hand-wave into an explanation.
Ionic bonding
Ionic bonds form between a metal and a non-metal. It is an electrostatic attraction that acts in all directions, holding every ion to every oppositely charged neighbour throughout a giant lattice.
Ionic bonding — the electrostatic attraction between oppositely charged ions in a lattice.
The ions form by electron transfer. At A-Level the configurations are written in full: sodium (1s22s22p63s1) loses its 3s electron to become Na+ (1s22s22p6), and chlorine (1s22s22p63s23p5) gains it to become Cl− (1s22s22p63s23p6) — configuration rules are in Atomic Structure.
The ions then pack into a giant ionic lattice — a regular, repeating arrangement of alternating positive and negative ions, extending in three dimensions. In sodium chloride each Na+ touches six Cl− and each Cl− touches six Na+ (6:6 co-ordination). There is no such thing as “a molecule of NaCl”: the formula just states the ratio of ions in the lattice.
Predicting charges & building formulas
The periodic table hands you the charge on a simple ion: metals lose their outer electrons, non-metals gain enough to fill the outer sub-shells.
Charge = the group number
Charge = the group number − 8
No pattern — recall these five cold. Ammonium is the only positive one.
Ammonium previews the next section — its four N–H bonds are covalent, one of them dative.
Step 1 — write the two ions:
Al3+ and SO42−
Step 2 — balance the charges to zero (lowest common multiple of 3 and 2 is 6):
2 × (3+) = 6+ and 3 × (2−) = 6−
Step 3 — write the formula, brackets round the repeated compound ion:
Al2(SO4)3
🧪 Exam-style questions
Which statement about inorganic ionic compounds is always correct? Tick (✓) one box.
Source: AQA A-Level Chemistry past papers (2020–2024).
Formula construction and ionic equations are drilled properly in Amount of Substance — formulae & equations, with extra practice in the Bridge the Gap formula drills. GCSE groundwork: C2 Part 2 — ionic bonding & lattices.
A sodium atom (1s22s22p63s1) and a sodium ion (1s22s22p6) are different species with different sizes, different configurations and completely different chemistry. In any sentence about ionic substances, the mark-bearing word is ions: ionic lattices contain no atoms and no molecules. Their sizes have different names, too — the atomic radius describes the atom, the ionic radius the ion (a positive ion is smaller than its atom, a negative ion larger). Keep the two apart in any size or lattice-strength answer.
- “Electrostatic attraction between oppositely charged ions” is the full phrase — recent AQA examiner reports flag answers that say atoms or molecules where ions are required.
- Ionic bonding is the attraction, not the transfer. Electron transfer only makes the ions; the bonding is what holds them together afterwards.
- Formulas balance charge, and compound ions travel whole — Al2(SO4)3, never Al2S3O12.
The strength of an ionic lattice — and with it the melting point — grows with the charge on the ions and shrinks as the ions get larger (the charges sit further apart). Putting an energy value on that attraction is lattice enthalpy, a Year 2 story.
Covalent & dative covalent bonds
Where ionic bonding transfers electrons, covalent bonding shares them between non-metals— and the shared pair is attracted to both nuclei at once. That attraction, pulling two nuclei toward the same pair of electrons, is the bond.
A single covalent bond — the strong electrostatic attraction between the positively charged nuclei of two atoms and the shared pair of electrons between them. Multiple bonds contain multiple shared pairs: a double bond two pairs, a triple bond three.
In displayed formulas a covalent bond is drawn as a line — one line per shared pair. So O=O carries two shared pairs and N≡N three, which is why nitrogen is so unreactive.
The dative covalent (co-ordinate) bond
A co-ordinate (dative covalent) bond — a shared pair of electrons in which both electrons are supplied by one atom.
Two requirements, always: the donor needs a lone pair, and the acceptor needs an empty orbital to receive it. The bond is drawn as an arrow pointing from the lone pair to the acceptor. The classic case is ammonia protonating:
NH3 + H+ → NH4+
H+ has no electrons at all, so nitrogen’s lone pair supplies both. The same move makes the oxonium ion, H3O+, when water grabs a proton — the ion behind every acid you will meet. And ammonia’s lone pair can plug into the empty orbital on boron in BF3 (boron has only six electrons around it), forming H3N→BF3. Aluminium chloride pulls the same trick with itself, pairing up into Al2Cl6 through two dative bonds.
Yes — the new N–H bond in NH4+ is a dative bond, so it is drawn as an arrow: nitrogen supplied both of its electrons (the incoming H+ brought none). But a dative bond is a normal covalent bond — same length, same strength as any other bond between the same two atoms. The arrow records the origin of the electrons, nothing more. Once the ion has formed, all four N–H bonds are identical and you cannot tell which one was dative — which is why the arrow is only a bookkeeping mark, not a different kind of bond.
- The arrow points from the donor. It starts at the lone pair and ends at the acceptor — a reversed arrow describes a different (wrong) electron flow.
- Draw the lone pairs. Recent AQA examiner reports flag diagrams that omit them — and a dative-bond explanation without a named lone pair and empty orbital is incomplete.
- “Dative” describes a bond inside a particle. It is not a force between molecules — the contrast with hydrogen bonding is drawn properly in Forces between molecules.
🧪 Exam-style questions
Which compound contains a co-ordinate (dative) bond? Tick (✓) one box.
Which molecule can accept an electron pair during the formation of a co-ordinate bond? Tick (✓) one box.
Which reaction involves the formation of a dative covalent bond? Tick (✓) one box.
Source: AQA A-Level Chemistry past papers (2020–2024).
GCSE groundwork: C2 Part 3 — covalent substances. Dative bonding returns as the way ligands attach in transition-metal complexes.
Metallic bonding
A metal is a lattice of positive ions — not atoms — sitting in a sea of electrons that no longer belong to any one ion. Each metal atom releases its outer-shell electrons into the shared pool; the electrons are delocalised, free to move throughout the whole structure.
Metallic bonding — the electrostatic attraction between delocalised electrons and positive ions arranged in a lattice.
How strong the bonding is comes down to three linked factors — compare sodium (melts at 98 °C) with magnesium (650 °C):
- Charge on the ion: Mg2+ attracts the electron sea more strongly than Na+.
- Number of delocalised electrons: magnesium contributes two per ion, sodium one.
- Size of the ion: Mg2+ is smaller, so the electrons sit closer to the nuclear charge.
- Positive ions, delocalised electrons — a metal lattice contains no neutral atoms and no anions. “Attraction between atoms and electrons” scores nothing.
- Comparisons need both sides. To explain magnesium’s higher melting point, state the charge, electron count and size for both metals — one property asserted for one metal is half an answer.
- Metals are not ionically bonded. There are no oppositely charged ions present — the attraction is ion-to-electron.
🧪 Exam-style questions
Which element has the lowest melting point? Tick (✓) one box.
Describe the bonding in magnesium.
Show answer
A lattice of magnesium Mg2+ ions surrounded by a sea of delocalised electrons (the outer-shell electrons). 1 mark
The bonding is the electrostatic attraction between the positive ions and the delocalised electrons. 1 mark
Source: AQA A-Level Chemistry past papers (2020–2024).
GCSE groundwork: C2 Part 4 — metallic bonding. The Na → Mg → Al melting-point climb across Period 3 runs on exactly these factors.
The four crystal structures
Every crystalline solid on this course is one of four types — ionic, metallic, macromolecular (giant covalent) or molecular — and six named examples carry all of them: sodium chloride, magnesium, diamond, graphite, ice and iodine. Identify the type first; every property answer follows from it.
Ionic & metallic crystals
Sodium chloride is the ionic example — the alternating 6:6 lattice drawn above, held by electrostatic attraction between oppositely charged ions. Magnesium is the metallic example — the ion lattice in its electron sea, also drawn above. Both are giant structures: strong attractions repeated throughout, so both are high-melting solids.
Macromolecular (giant covalent) crystals
Diamond: every carbon is covalently bonded to four others in a tetrahedral network that continues in all directions. No discrete molecules, no spare electrons — just covalent bonds from edge to edge of the crystal, which is why diamond is exceptionally hard, does not conduct, and does not melt below roughly 3,500 °C.
Graphite: every carbon bonds to three others, building flat layers of hexagons. The fourth outer electron from each carbon is delocalised between the layers — so graphite conducts electricity. The layers themselves are held to each other only by weak forces (induced dipole–dipole — named properly in Forces between molecules), so they slide: graphite is soft enough to write with.
Molecular crystals
A molecular crystal is a lattice of molecules. Inside each molecule the atoms are held by strong covalent bonds; between the molecules there are only weak intermolecular forces — and it is those that break on melting.
Iodine: a regular arrangement of I2 molecules. The I–I covalent bond inside each molecule is strong and survives melting untouched; the induced dipole–dipole forces between molecules are weak, so iodine melts at just 114 °C and sublimes readily to a purple vapour of intact I2 molecules.
Ice is also molecular, with a twist: hydrogen bonds (the strongest of the intermolecular forces, defined in Forces between molecules) hold the H2O molecules in an open lattice with large gaps. When ice melts, that open arrangement collapses and the molecules move closer together — which is why ice is less dense than liquid water and floats.
“Iodine” names two things: the molecule I2, and the crystal — a lattice of those molecules. Properties of the crystal (melting, subliming) are about the forces between molecules; the covalent bond within each molecule is a separate, far stronger thing. Keeping the two apart is most of this topic’s marks.
AQA asks you to “draw diagrams to represent these structures involving specified numbers of particles” (AQA 3.1.3.4), so a diagram question gives you a number of particles to show. To make it score:
- Draw the number asked for — and in the right arrangement: alternating ions for a lattice, a bonded network for giant covalent.
- Label the actual particles — Na+ and Cl− with their charges, not “Na” and “Cl”.
- Show a giant structure continuing — leave bonds dangling at the edges of a diamond or graphite fragment; a closed cage reads as a single molecule.
- Show a molecular crystal as whole molecules — discrete I2 units in a regular pattern, never a continuous network.
🧪 Exam-style questions
In which substance do covalent bonds break when it melts? Tick (✓) one box.
Which substance contains delocalised electrons? Tick (✓) one box.
Source: AQA A-Level Chemistry past papers (2020–2024).
GCSE groundwork: C2 Part 3 — giant covalent structures and C2 Part 1 — states of matter. Period 3 replays all four structure types across a single row of the periodic table.
Melting points, conductivity & changes of state
Property explanations on this topic are a fixed routine. Run it in order, every time:
- Name the structure type — ionic / metallic / macromolecular / molecular.
- Name the particles — ions, ions + delocalised electrons, atoms, or molecules.
- Name what is overcome (melting) or what carries charge (conductivity).
- Compare its strength — or the carrier’s mobility.
- Conclude — energy needed, so melting point; carriers mobile or not, so conducts or not.
| Crystal type | Particles | Overcome on melting | Melting point | Conducts as solid? | Conducts molten / in solution? | Examples |
|---|---|---|---|---|---|---|
| Ionic | Oppositely charged ions | Electrostatic attraction between ions | High (NaCl 801 °C) | No — ions fixed in the lattice | Yes — ions free to move | Sodium chloride |
| Metallic | Positive ions + delocalised electrons | Attraction between ions and electrons | Generally high (Mg 650 °C) | Yes — delocalised electrons move | Yes (molten) | Magnesium |
| Macromolecular | Atoms, covalently bonded network | Covalent bonds (many, strong) | Very high (diamond > 3,500 °C) | No — except graphite’s delocalised electrons | — (does not dissolve; barely melts) | Diamond, graphite |
| Molecular | Molecules | Weak intermolecular forces (bonds inside molecules untouched) | Low (ice 0 °C, iodine 114 °C) | No — no charged particles at all | No | Ice, iodine |
Conduction needs mobile charged particles — that is the whole test. Metals and graphite carry current on delocalised electrons in any state; ionic compounds only once the lattice is broken up by melting or dissolving, freeing the ions (the electrons stay put); molecular substances have no charged particles to offer, solid or liquid.
Melting point and conductivity are the two properties you must be able to explain, but two others help you pin down a structure — and both trace straight back to the particles:
- Solubility in water: ionic compounds often dissolve, because polar water molecules pull the ions out of the lattice (attractions between the ions and the δ+/δ− of water). Molecular substances usually do not dissolve in water — unless they are polar or can hydrogen-bond with it (sugar and ethanol dissolve; iodine and hexane do not). Giant covalent solids and metals are insoluble. (“Polar” and “hydrogen-bond” are explained in Electronegativity & bond polarity and Forces between molecules below.)
- Mechanical behaviour: ionic crystals are hard but brittle — knock one layer along and ions of like charge line up, repel, and the crystal splits. Metals are malleable and ductile — layers of ions slide over one another while the delocalised electrons keep holding them together, so the metal changes shape without shattering. Among the giant covalent solids, diamond is exceptionally hard (a rigid 3-D network of covalent bonds) while graphite is soft and slippery (its layers slide).
Energy & changes of state
Melting and boiling cost energy because attractions between particles must be overcome — which attraction, the table above tells you. While a state change is happening, the energy supplied goes into separating particles rather than speeding them up, so the temperature holds constant at the melting or boiling point. The stronger the attraction, the more energy, the higher the temperature at which it happens. Iodine’s forces are so weak it can skip the liquid stage entirely and sublime.
Explain why sodium chloride melts at 801 °C but iodine melts at 114 °C.
Sodium chloride is an ionic lattice: melting it means overcoming the strong electrostatic attraction between oppositely charged ions, which acts throughout the giant structure, so a large amount of energy — and a high temperature — is needed. Iodine is a molecular crystal: melting it only means overcoming the weak induced dipole–dipole forces between I2 molecules, while the strong I–I covalent bonds inside the molecules are not broken. Far less energy is needed, so iodine melts at a much lower temperature.
Notice the routine underneath: structure type → particles → attraction overcome → strength comparison → energy conclusion. Written as full sentences, every move scores.
Three bench tests pin down a structure type: ease of melting (melts in a Bunsen flame, or not), electrical conductivity (test the solid, then molten or in solution), and solubility in water. A white solid that won’t melt and never conducts points macromolecular; one that dissolves and conducts only in solution points ionic; a low-melting non-conductor points molecular. This experiment is fair game in practical contexts.
- Melting a molecular substance never breaks covalent bonds. Ice melting breaks hydrogen bonds between H2O molecules — every O–H bond survives. “The covalent bonds break” is the classic zero-scorer on this topic.
- Giant structures are the opposite: melting diamond or NaCl really does break the network’s bonds or the lattice’s ionic attraction — the answer depends on structure type, which is why the routine names it first.
- “Contains electrons” explains nothing — everything contains electrons. Conduction answers name the mobile carrier: delocalised electrons, or ions free to move once molten or dissolved.
🧪 Exam-style questions
Which is not responsible for conducting electricity? Tick (✓) one box.
The table shows the melting points of three substances: sodium chloride 1074 K, chlorine 172 K, hydrogen chloride 158 K. Explain why these melting points are so different. Refer to the structure of, and bonding in, each substance.
Show answer
Sodium chloride — a giant ionic lattice; melting overcomes the strong electrostatic attraction between the oppositely charged Na+ and Cl− ions throughout the lattice, so its melting point is by far the highest. 2 marks
Chlorine and hydrogen chloride are both simple molecular — melting overcomes only weak intermolecular forces, so both melt far lower than NaCl. 2 marks
Cl2 melts a little higher than HCl even though HCl is polar: Cl2 has more electrons, so its induced dipole–dipole (van der Waals) forces are stronger than the permanent dipole–dipole forces in HCl. 2 marks
The discriminator is that last point — more electrons (Cl2) can outweigh a permanent dipole (HCl).
Source: AQA A-Level Chemistry past papers (2020–2024).
GCSE groundwork: C2 Part 5 — does it conduct?. Putting numbers on bond making and breaking is Energetics (3.1.4).
Shapes of molecules & ions
Electron pairs are charge clouds that repel one another, so the pairs in an atom’s outer shell arrange themselves as far apart as possible to minimise repulsion. That single idea — electron pair repulsion — predicts the shape of any molecule or ion on this course. One refinement completes it: not all repulsions are equal.
lone pair–lone pair > lone pair–bond pair > bond pair–bond pair
A lone pair sits closer to the central atom than a bonding pair, so it repels harder. Each lone pair squeezes the remaining bond angles by about 2.5°.
Counting the pairs
Any species, familiar or not, yields to the same five steps:
- Outer-shell electrons on the central atom = its group number.
- Add one per atom bonded (each brings one electron to its bond).
- Adjust for charge: add one per negative charge, subtract one per positive.
- Divide by two → total electron pairs. Total − bonded atoms → lone pairs.
- Name the shape and state the angle from the pair combination (chart below).
Multiple bonds: treat a double bond as one charge cloud — CO2 has two clouds around carbon, so it is linear.
Linear
180° · BeCl2, CO2
Trigonal planar
120° · BF3
Bent (v-shaped)
~119° · SO2
Tetrahedral
109.5° · CH4
Trigonal pyramidal
107° · NH3
Bent (v-shaped)
104.5° · H2O
Trigonal bipyramidal
120° & 90° · PCl5
See-saw
~102° & ~89° · SF4
T-shaped
~89° · ClF3
Linear
180° · XeF2
Octahedral
90° · SF6
Square pyramidal
~89° · BrF5
Square planar
90° · XeF4
Two reading rules for the chart. First, the name describes where the atoms are, not where the electron pairs are — ICl2− has five pairs arranged around a trigonal bipyramid, but with three of them lone, the three atoms sit in a line: linear. Second, in five-pair arrangements lone pairs always take the roomier equatorial positions — which is why XeF2 is linear and ClF3 T-shaped.
Explaining a shape for full marks
Shape explanations are three sentences, always the same three:
1. State the pairs: “There are X bonding pairs and Y lone pairs of electrons around the central atom.”
2. State the principle: “Electron pairs repel and arrange themselves as far apart as possible.”
3. If lone pairs are present, rank the repulsion: “Lone pairs repel more than bonding pairs, so the bond angle is reduced from … to …”
For ammonia: “There are 3 bonding pairs and 1 lone pair around the nitrogen atom. The electron pairs repel and move as far apart as possible. The lone pair repels more than the bonding pairs, so the bond angle is reduced from 109.5° to 107°.”
Step 1 — ICl2−: count the electrons on the central iodine:
7 (Group 7) + 2 (two bonded Cl) + 1 (one negative charge) = 10 electrons
Step 2 — convert to pairs and lone pairs:
10 ÷ 2 = 5 pairs; 5 − 2 bonded atoms = 3 lone pairs
Step 3 — name the shape from 2 bonding + 3 lone:
lone pairs take the equatorial positions → the atoms lie in a line → linear, 180°
Step 4 — PCl4+, same routine:
5 + 4 − 1 = 8 electrons → 4 pairs, 0 lone → tetrahedral, 109.5°
The charge step is where unfamiliar ions are won or lost: add electrons for negative charges, remove them for positive.
Interactive — deduce the shape, one step at a time
Work out the shape of .
Step 1 — how many outer-shell electrons does the central atom () have? (= its group number)
Step 2 — add one electron per bonded atom, adjust for charge, then ÷ 2. How many charge clouds (bonding regions + lone pairs)? Count each double bond as one cloud.
Step 3 — split those into bonding pairs and lone pairs (lone pairs = clouds − bonded atoms).
Step 4 — name the shape (the name describes where the atoms are).
Step 5 — give the bond angle.
- Name the shape exactly, and draw every lone pair. Recent AQA examiner reports flag close-but-unnamed drawings and diagrams with lone pairs missing — the name and the lone pairs each carry marks.
- “They repel” is not an explanation. When lone pairs are involved, state which repulsion is greater — recent reports single out bond-angle answers of 120° or 180° given for molecules like water, from students who never counted the pairs.
- Count before you name. The routine exists because intuition fails on ions — run the five steps even when the species looks familiar.
🧪 Exam-style questions
Which species has a square planar shape? Tick (✓) one box.
Which species has a lone pair of electrons on the central atom? Tick (✓) one box.
SF6 and SF3+ have different shapes and different bond angles. Deduce the shape of each species, state the bond angle in each, and justify the angles by referring to electron pairs.
Show answer
SF6: 6 bonding pairs, 0 lone pairs — octahedral, bond angle 90°. The six pairs repel equally and spread as far apart as possible. 3 marks
SF3+: 3 bonding pairs and 1 lone pair — trigonal pyramidal, bond angle about 107° (accept 103–107°). The lone pair repels more than the bonding pairs, pushing the S–F bonds closer together, below the tetrahedral 109.5°. 3 marks
Source: AQA A-Level Chemistry past papers (2020–2024).
Outer-shell counting works because of the configuration rules in Atomic Structure. Shapes return with complex ions and wedge–dash drawing pays off across organic chemistry.
Electronegativity & bond polarity
A covalent bond is a shared pair — but the sharing is rarely equal. How hard each atom pulls on the pair is its electronegativity.
Electronegativity — the power of an atom to attract the pair of electrons in a covalent bond. Measured on the Pauling scale; no units.
Electronegativity increases across a period (nuclear charge rises, radius shrinks, shielding stays the same) and decreases down a group (radius and shielding grow) — the same three factors that control ionisation energy in Atomic Structure. So the most electronegative atoms cluster top-right: fluorine (4.0) leads, then oxygen (3.5), with nitrogen and chlorine (3.0) close behind. Carbon sits at 2.5 and hydrogen 2.1.
Polar bonds
Bond two atoms of different electronegativity and the shared pair sits nearer the stronger puller: the electron distribution is unsymmetrical. The bond is polar — the pulling atom carries a partial negative charge (δ−), its partner a partial positive (δ+). In H–Cl, chlorine (3.0) out-pulls hydrogen (2.1): Hδ+–Clδ−. The bigger the electronegativity difference, the more polar the bond.
Identical atoms share equally (a pure covalent bond, Cl2); a modest difference skews the pair (polar covalent, HCl); a large difference transfers it outright (ionic, NaCl). The three bond types of this page are the two ends and the middle of one continuous scale.
Polar bonds vs polar molecules
A polar bond does not guarantee a polar molecule. Each polar bond contributes a small dipole with a direction; whether the molecule as a whole has a permanent dipole depends on whether those bond dipoles cancel. Symmetry decides — which is why this section needed Shapes first. CO2 (linear), BF3 (trigonal planar) and CCl4 (tetrahedral) all contain polar bonds arranged so symmetrically that the dipoles cancel exactly: no permanent dipole. H2O, NH3 and CHCl3 are asymmetrical: the dipoles add up, and the molecule is polar.
Polar bond: a property of one bond — two atoms, unequal pull, δ+/δ−. Polar molecule (permanent dipole): a property of the whole molecule — it needs polar bonds and an asymmetrical shape, so the bond dipoles do not cancel. Every polarity answer states both halves.
Both molecules contain polar C–Cl bonds. Explain why CHCl3 has a permanent dipole but CCl4 does not.
In CCl4 the four C–Cl bond dipoles are identical and arranged tetrahedrally, so they are symmetrical and cancel: no permanent dipole. Replacing one Cl with H breaks that symmetry — the three remaining C–Cl dipoles no longer cancel against the C–H bond, so CHCl3 has a net dipole and the molecule is polar.
- A molecule is never polar “because it hydrogen-bonds”. It is polar because it has polar bonds and an asymmetrical shape — recent AQA examiner reports flag exactly this confusion. (Hydrogen bonding is a consequence of polarity, next section.)
- Anchor every comparison in electronegativity: “chlorine is more electronegative than hydrogen” — recent reports note vague polarity language as a recurring weakness.
- δ symbols sit on atoms, not floating mid-bond — and the arrow points toward δ−.
🧪 Exam-style questions
Which bond has the most unsymmetrical electron distribution? Tick (✓) one box.
Which species contains bonds that have different polarities? Tick (✓) one box.
Which molecule does not have a permanent dipole? Tick (✓) one box.
Source: AQA A-Level Chemistry past papers (2020–2024).
Bond polarity decides where attacking species strike in organic mechanisms.
Forces between molecules
Molecular substances are held together twice over. Within each molecule, the atoms are joined by strong covalent bonds; these are the intramolecular forces (intra‑ means within). Between one molecule and the next act much weaker attractions; these are the intermolecular forces (inter‑ means between, as in an international flight, which travels between nations). Three intermolecular forces cover everything on this course, and telling them apart — and apart from the covalent bonds — is where the marks live.
A covalent bond is a shared pair binding atoms into a molecule — typically hundreds of kJ mol−1 to break. Intermolecular forces are attractions between whole molecules — weaker by an order of magnitude or more. Melting and boiling a molecular substance overcomes only the forces between molecules; the molecules themselves arrive in the gas intact.
Strip away the names and every attraction in this topic — ionic, covalent and metallic bonding, and all three intermolecular forces — is at heart the same thing: an electrostatic attraction between opposite charges. What changes is the strength. We give them separate names and rank them — the three bonds strong, then hydrogen bonding, then permanent dipole–dipole, then induced dipole–dipole progressively weaker — only so we can compare them and explain properties. Underneath, it is one idea at different sizes.
That is also why hydrogen bonding is awkwardly named. Despite the word “bond”, it is an intermolecular force, not a chemical bond — no pair of electrons is shared and no new substance forms. It picked up “bond” historically because it is by far the strongest intermolecular force, strong enough to behave a little like one; but in an answer it belongs with the forces between molecules, never with covalent bonds.
Induced dipole–dipole forces
AQA’s name is induced dipole–dipole forces; the same force is called van der Waals, dispersion or London forces elsewhere. These notes use induced dipole–dipole throughout.
Electrons are in constant motion, so at any instant a molecule’s electron cloud can sit lopsided: an instantaneous dipole. That flicker of charge induces an opposite dipole in the neighbouring molecule, and the two attract. The dipoles come and go; the attraction, averaged over time, persists. Because it needs nothing but electrons, this force operates between all molecules, polar or not — and it strengthens as the number of electrons grows (bigger, floppier clouds polarise more easily). More contact between molecules also helps — one reason unbranched chains outboil their branched isomers.
Permanent dipole–dipole forces
Polar molecules carry their dipoles permanently, so the δ+ of one attracts the δ− of the next: permanent dipole–dipole forces. They act in addition to induced dipole–dipole forces, never instead of them — so two substances with similar electron counts usually see the polar one boil higher.
Propanone (CH3COCH3, 32 electrons) boils at 56 °C; butane (C4H10, 34 electrons) boils at −0.5 °C. Explain the difference.
The two molecules have almost the same number of electrons, so their induced dipole–dipole forces are about equal. Butane is non-polar, so induced dipole–dipole is all it has. Propanone is polar (the C=O bond gives it a permanent dipole), so it has permanent dipole–dipole forces on top of its induced dipole–dipole forces. More energy is needed to overcome the intermolecular forces of attraction between propanone’s molecules, so it boils far higher — the fair-test way to isolate the effect of the permanent dipole is to compare molecules of similar electron count.
Hydrogen bonding
The strongest intermolecular force is a special case with strict entry requirements. Hydrogen bonded directly to nitrogen, oxygen or fluorine — the three small, intensely electronegative atoms — is stripped almost bare of electron density: an unusually large δ+ with no inner shells to shield it. A lone pair on the N, O or F of another molecule is attracted straight to it.
A hydrogen bond is the attraction between a lone pair on N, O or F and a δ+ hydrogen atom covalently bonded to N, O or F in another molecule. No N–H, O–H or F–H bond in the molecule — no hydrogen bonding, however polar the molecule is (HCl has none).
That shared starting point is exactly why they get swapped. A dative bond shares the lone pair into an empty orbital: a full covalent bond within one particle, drawn as an arrow. A hydrogen bond is an attraction between separate molecules: far weaker, nothing shared into an orbital, drawn as a dashed line. And neither is the same as permanent dipole–dipole: hydrogen bonding needs H on N, O or F specifically — polar HCl manages only permanent dipole–dipole.
What the forces do to boiling points
Boiling separates molecules, so boiling points read out intermolecular force strength directly. Down a series of similar molecules, more electrons → stronger induced dipole–dipole forces → higher boiling point: the halogens go gas → gas → liquid → solid down the group, and HCl → HBr → HI boils progressively higher even though the bonds get less polar — strong evidence that induced dipole–dipole usually outweighs permanent dipole–dipole.
Then three hydrides jump off the trend entirely. NH3, H2O and HF boil far above what their electron counts predict, because each hydrogen-bonds. Water outdoes them all: with two lone pairs and two δ+ hydrogens per molecule, it averages around two hydrogen bonds per molecule where HF manages about one — hence a 100 °C boiling point on just ten electrons.
The same hydrogen bonds explain ice: held at full stretch in the open lattice drawn earlier, water’s molecules occupy more space frozen than liquid — so ice floats. Low density in a solid is rare, and hydrogen bonding is the reason water manages it.
Longer chain (a homologous series, e.g. butane → pentane → hexane): each extra CH2 adds more electrons, so the induced dipole–dipole forces get stronger and the boiling point rises. Here “more electrons” is the reason.
Same formula, different shape (isomers, e.g. pentane vs 2-methylbutane): the molecules have the same number of electrons, so electron count and “chain length” cannot be the reason for any difference — it comes down to surface contact: the more branched molecule touches its neighbours over a smaller area, so its induced dipole–dipole forces are weaker and it boils lower.
The classic exam trap gives you two isomers of the same molecular formula and rewards surface contact / branching — answers that reach for “more electrons” or “a longer chain” score nothing, because both molecules have identical electron counts.
Explain why HF boils at 20 °C while HCl boils at −85 °C, yet boiling points rise again from HCl to HI.
HF molecules form strong hydrogen bonds between molecules. HCl cannot hydrogen-bond, so its molecules are held only by permanent dipole–dipole and induced dipole–dipole forces, which need far less energy to overcome. From HCl to HI the molecules gain electrons, so the induced dipole–dipole forces strengthen and the boiling points increase.
Quick-fire — name the strongest force present
Every molecule has induced dipole–dipole forces. What is the strongest force acting between molecules of ?
Run a fine jet of liquid from a burette past a charged rod: a polar liquid (water) bends visibly toward it, while a non-polar one (hexane) falls straight. Relative deflection is a quick probe for permanent dipoles — a neat practical link for this topic.
- “Hydrogen bonds between oxygen and hydrogen” is too vague to score. Write the full explanation the mark scheme wants: a lone pair on the O (or N/F) of one molecule attracts the δ+ H bonded to O (or N/F) in another.
- Hydrogen bond ≠ dative bond. Recent reports note each being named when the other is asked for — if electrons are shared into an empty orbital it is dative; if two molecules attract, it is hydrogen bonding.
- Boiling breaks intermolecular forces only. Steam is made of intact H2O molecules — no O–H bond breaks at 100 °C.
- No molecule is force-free. Every molecule has induced dipole–dipole forces — “no intermolecular forces” is never a correct description.
- Name the specific force, in full. At GCSE “weak intermolecular forces” earned the mark; at A-Level it no longer does — you must name the type present (induced dipole–dipole, permanent dipole–dipole or hydrogen bonding) and write it out in words. The abbreviation “IMF” is not accepted in an exam answer.
🧪 Exam-style questions
Which compound has hydrogen bonding? Tick (✓) one box.
Which of these alkanes has the highest boiling point? Tick (✓) one box.
Dichloromethane (CH2Cl2, polar, b.p. 40 °C), tetrachloromethane (CCl4, non-polar, b.p. 77 °C) and propan-1-ol (polar, b.p. 97 °C). (a) State why the C–Cl bonds are polar. (b) Suggest why CCl4 molecules are non-polar. (c) Explain why CCl4 boils higher than CH2Cl2. (d) Describe the hydrogen bonding in propan-1-ol.
Show answer
(a) Chlorine is more electronegative than carbon, so the shared electrons are pulled unequally — the bond is polar. 1 mark
(b) CCl4 is symmetrical (tetrahedral), so its four C–Cl bond dipoles cancel. 1 mark
(c) CCl4 has more electrons than CH2Cl2, so its induced dipole–dipole (van der Waals) forces are stronger — strong enough to outweigh the permanent + induced dipole–dipole forces in the smaller CH2Cl2 — so more energy is needed to separate its molecules. 2 marks
(d) A lone pair on the O of one propan-1-ol molecule attracts the δ+ H of the O–H group on a neighbouring molecule. 2 marks
Source: AQA A-Level Chemistry past papers (2020–2024).
GCSE groundwork: C2 Part 3 — simple molecules. Hydrogen bonding returns in alcohols and induced dipole–dipole trends in the boiling points of alkanes.
Capstone: name the attraction
Every property question on this topic is the same three decisions, made before any explaining starts: what are the particles, what attraction acts between them, and what happens to it. Get those three right and the five-move explanation writes itself; get the first one wrong and nothing after it can score.
The generator below drills the first two moves: classify a substance into its structure type, then name the attraction that must be overcome to melt it — the step where the covalent-bonds-versus-intermolecular-forces trap catches people out.
| Decision chain | Sodium chloride | Diamond | Ice |
|---|---|---|---|
| Particles | Na+ and Cl− ions | Carbon atoms, covalently networked | H2O molecules |
| Attraction between them | Electrostatic attraction between oppositely charged ions | Covalent bonds throughout | Hydrogen bonds (+ induced dipole–dipole) |
| On melting | Ionic attraction overcome — lots of energy, 801 °C | Covalent bonds broken — huge energy, > 3,500 °C | Hydrogen bonds overcome, O–H bonds untouched — 0 °C |
| Conducts? | Only molten/dissolved — ions mobile | No — no mobile charge (graphite’s delocalised electrons are the exception) | No charged particles |
Capstone — classify the substance, then name the attraction
1 — classify. Which structure type is ?
2 — name the attraction. What has to be overcome to melt it?
3 — name the force. Which is the strongest intermolecular force between its molecules?
- Ionic bonding — electrostatic attraction between oppositely charged ions in a lattice; NaCl is 6:6 co-ordinated; no molecules anywhere in it.
- Ion charges — from table position: +1/+2/+3 (Groups 1–3), −3/−2/−1 (Groups 5–7); d-block charges named with Roman numerals; compound ions recalled cold: SO42−, OH−, NO3−, CO32−, NH4+.
- Formulas — charges cancel to zero; brackets round repeated compound ions: Al2(SO4)3.
- Covalent bond — a shared pair attracted to both nuclei; one line per pair; double and triple bonds = two and three pairs.
- Dative (co-ordinate) bond — both electrons from one atom: lone pair → empty orbital, drawn as an arrow from the donor; once formed, identical to any covalent bond (NH4+, H3O+, H3N→BF3).
- Metallic bonding — electrostatic attraction between delocalised electrons and positive ions in a lattice; stronger with higher charge, more electrons per ion, smaller ions (Mg 650 °C vs Na 98 °C).
- Four crystal types — ionic (NaCl), metallic (Mg), macromolecular (diamond: 4 bonds per C; graphite: 3 + a delocalised electron, sliding layers), molecular (ice, iodine: lattices of molecules).
- Melting — overcome whatever holds the particles: ionic attraction, ion–electron attraction, covalent network, or (molecular) weak intermolecular forces only — covalent bonds inside molecules never break on melting.
- Conductivity — needs mobile charged particles: delocalised electrons (metals, graphite) or ions freed by melting/dissolving; molecular substances never conduct.
- Changes of state — energy overcomes attractions between particles; temperature plateaus during the change; stronger attraction → higher melting/boiling point; weak-forces solids like iodine can sublime.
- Shapes routine — group number + bonded atoms ± charge, ÷2 → pairs; pairs − bonded atoms → lone pairs; double bond = one charge cloud; name describes atom positions.
- Repulsion hierarchy — lp–lp > lp–bp > bp–bp; each lone pair closes the angle ~2.5°; angles to quote: 180°, 120°, 109.5°, 107°, 104.5°, 120° & 90°, 90°.
- Electronegativity — the power of an atom to attract the pair of electrons in a covalent bond; increases across a period, decreases down a group; F highest.
- Polar molecule — polar bonds and an asymmetrical shape (dipoles don’t cancel); symmetric CO2, BF3, CCl4 have polar bonds but no permanent dipole.
- Three intermolecular forces — induced dipole–dipole (in every molecule; grows with electron count; also called van der Waals, London or dispersion forces), permanent dipole–dipole (polar molecules, on top), hydrogen bonding (lone pair on N/O/F ↔ δ+ H bonded to N/O/F on another molecule; dashed line, ~180°). The forces stack — a hydrogen-bonder also has the other two; the strongest present sets the boiling point.
- Chain length vs branching — more carbons down a series means more electrons and stronger induced dipole–dipole (higher b.p.); but between isomers of the same formula the electron count is equal, so any difference is down to surface contact — straight chains touch over more area than branched ones. Quoting “more electrons” for isomers is a common error.
- Hydrogen-bonding consequences — NH3, H2O, HF boil far above their group trends (water ~2 H-bonds per molecule); ice’s open H-bonded lattice makes it less dense than liquid water.