At GCSE
At GCSE you draw dot-and-cross diagrams for ionic, covalent and metallic bonding and link structure (giant ionic, simple molecular, giant covalent, metallic) to properties such as melting point, conductivity and solubility.
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A-Level Chemistry Chemical Bonding
This is a chemical bonding revision resource for A-Level Chemistry, written for UK students and matched to the bonding content examined by AQA (7405), Edexcel and OCR. It also works as a chemical bonding GCSE refresher, since the ionic, covalent and metallic ideas carry straight through.
Work down the page in order or jump to a section. Every part links bonding and structure to the properties examiners ask you to explain: melting point, conductivity, solubility and the shapes and polarity of molecules.
What is ionic bonding?
Ionic bonding is the strong electrostatic attraction between oppositely charged ions. A metal atom loses one or more electrons to reach a stable electron arrangement and a non-metal atom gains them, so you end up with cations and anions held in a giant ionic lattice rather than in separate molecules.
Marks usually come from the word electrostatic and from describing the lattice. Saying that sodium chloride is made of NaCl molecules loses credit: the formula is simply the smallest whole-number ratio of ions in the lattice.
- Lattice enthalpy and melting point rise as ionic charge rises and ionic radius falls, so MgO melts far above NaCl.
- Solid ionic compounds do not conduct: the ions are fixed. Molten or aqueous they do, because the ions are mobile.
- Many ionic solids dissolve in water because polar water molecules attract and surround the separated ions.
- Self-test: explain in electrostatic terms why calcium oxide has a higher melting point than potassium chloride.
What is covalent bonding?
A covalent bond is a shared pair of electrons attracted to the nuclei of both bonded atoms. Multiple bonds share two or three pairs, and shorter bonds are generally stronger, which is why a C=O bond needs more energy to break than a C-O bond.
Distinguish simple molecular substances from giant covalent ones. In carbon dioxide or iodine the covalent bonds inside the molecules are strong, but only weak forces act between molecules, so melting points are low. In diamond or silicon dioxide the whole structure is one covalent network, so melting requires breaking strong bonds throughout.
- Graphite conducts because each carbon bonds to three others, leaving one delocalised electron per atom free to move along the layers.
- Diamond does not conduct: every outer electron is localised in a bond.
- Self-test: why does iodine sublime easily even though the I-I bond is strong?
What is a dative covalent bond?
A dative covalent (coordinate) bond is a covalent bond in which both shared electrons come from the same atom, normally from a lone pair. Once formed it behaves identically to any other covalent bond of the same type, with the same length and strength.
Show it with an arrow pointing from the lone-pair donor to the acceptor. Standard examples are the ammonium ion NH4+, where the nitrogen lone pair bonds to H+, and the aluminium chloride dimer Al2Cl6, where a chlorine lone pair donates into an electron-deficient aluminium atom. Ligands bonding to a transition metal ion do the same thing.
- All four N-H bonds in NH4+ are equivalent, and the ion is tetrahedral with bond angles of about 109.5 degrees.
- Self-test: state which species donates the electron pair when ammonia reacts with a hydrogen ion.
Metallic bonding
In a metal, positive ions sit in a regular lattice within a sea of delocalised outer-shell electrons. The bonding is the attraction between those cations and the delocalised electrons, and it acts in every direction through the structure.
Strength grows with the charge on the ion, the number of delocalised electrons per atom and a smaller ionic radius, so magnesium is much harder to melt than sodium. Metals conduct because the electrons move freely, and they are malleable because layers of ions can slide without the bonding being broken.
How does electronegativity affect bond polarity?
Electronegativity is the ability of an atom to attract the pair of electrons in a covalent bond towards itself. It increases across a period and decreases down a group, because a larger nuclear charge acting over a smaller atomic radius with similar shielding pulls bonding electrons more strongly.
When bonded atoms differ in electronegativity the electron density shifts towards the more electronegative atom, giving a permanent dipole and partial charges written as delta plus and delta minus. A large enough difference makes the bonding essentially ionic, so bonding sits on a continuum rather than in two separate boxes.
- A polar bond does not guarantee a polar molecule. In CO2 and CCl4 the individual dipoles cancel by symmetry, so there is no overall dipole.
- In water the two O-H dipoles do not cancel, because the bent shape leaves a net dipole.
- Self-test: explain why trichloromethane is polar but tetrachloromethane is not.
How do you determine molecular shape?
Use electron-pair repulsion theory. Count the bonding pairs and lone pairs around the central atom, then arrange them as far apart as possible, because electron pairs repel. Lone pairs repel more strongly than bonding pairs, so each lone pair reduces the bond angle by roughly 2.5 degrees.
Name the shape from the positions of the atoms, not the positions of all the electron pairs. Ammonia has four electron pairs arranged tetrahedrally but is described as trigonal pyramidal with a bond angle near 107 degrees.
- 2 bonding pairs, 0 lone: linear, 180 degrees, e.g. CO2.
- 3 bonding pairs, 0 lone: trigonal planar, 120 degrees, e.g. BF3.
- 4 bonding pairs, 0 lone: tetrahedral, 109.5 degrees, e.g. CH4.
- 3 bonding pairs, 1 lone: trigonal pyramidal, about 107 degrees, e.g. NH3.
- 2 bonding pairs, 2 lone: bent, about 104.5 degrees, e.g. H2O.
- 6 bonding pairs, 0 lone: octahedral, 90 degrees, e.g. SF6.
- Self-test: predict and justify the shape and bond angle of the hydroxonium ion H3O+.
What are intermolecular forces?
Intermolecular forces act between molecules and are far weaker than the covalent bonds inside them. Boiling a simple molecular substance overcomes these forces only, which is why the covalent bonds in water are untouched when a kettle boils.
Three types matter at A-Level, in increasing strength for molecules of similar size: London (induced dipole-dipole) forces, permanent dipole-dipole forces, and hydrogen bonding. London forces exist between all molecules and get stronger as the number of electrons rises, which explains the boiling point trend from fluorine to iodine.
- Hydrogen bonding requires hydrogen bonded directly to nitrogen, oxygen or fluorine, plus a lone pair on the electronegative atom of a neighbouring molecule.
- Water, ammonia and hydrogen fluoride boil anomalously high for their molecular mass because of hydrogen bonding.
- Ice is less dense than water because hydrogen bonds hold the molecules in an open lattice slightly further apart.
- Self-test: rank butane, propanone and propan-1-ol by boiling point and justify the order.
Structure and physical properties
Most exam questions on this topic really ask one thing: which forces are broken, and how strong are they? Identify the structure first, then name the force overcome on melting or boiling, then compare its strength.
- Giant ionic (NaCl, MgO): high melting point, conducts only when molten or aqueous, often water-soluble.
- Simple molecular (I2, CO2, H2O): low melting point, never conducts, solubility depends on polarity.
- Giant covalent (diamond, graphite, SiO2): very high melting point, insoluble; only graphite conducts.
- Metallic (Na, Mg, Al): high melting point, conducts as a solid and when molten, malleable and ductile.
At A-Level
At A-Level you add electronegativity, polarity and the three intermolecular forces (London, permanent dipole-dipole, hydrogen bonding). Expect VSEPR shape questions (bond angles, lone-pair repulsion) and explanations of anomalous boiling points such as water and HF.
Common exam mistakes
Saying covalent bonds break when a simple molecular substance boils.
Only the intermolecular forces are overcome. Name them precisely, for example hydrogen bonds between water molecules.
Confusing electron transfer with electron sharing.
Ionic bonding transfers electrons to form ions; covalent bonding shares a pair between two nuclei. Never describe an ionic compound as sharing electrons.
Assuming polar bonds always make a polar molecule.
Check the shape. Symmetrical arrangements such as CO2, BF3 and CCl4 cancel the individual dipoles, leaving no overall dipole.
Naming the electron-pair arrangement instead of the molecular shape.
Lone pairs count for repulsion but not for the shape name. Four pairs with one lone pair is trigonal pyramidal, not tetrahedral.
Writing that ionic compounds are made of molecules.
They form a giant lattice. The formula gives the simplest ratio of ions present.
Claiming hydrogen bonding occurs whenever a molecule contains hydrogen.
It needs H bonded directly to N, O or F, and a lone pair on a neighbouring N, O or F.
Explaining a melting point with vague wording such as stronger bonds.
State the structure, the force, and the reason it is stronger, for example a greater ionic charge and smaller radius giving stronger electrostatic attraction.
Chemical Bonding exam questions to test yourself
Original exam-style prompts. Answer out loud or in writing first, then open the answer to check your wording.
1. Define electronegativity and state two factors that make an atom more electronegative.+
The ability of an atom to attract the pair of electrons in a covalent bond. Greater nuclear charge and a smaller atomic radius with similar shielding both increase it.
2. Explain, in terms of structure and bonding, why magnesium oxide has a much higher melting point than sodium chloride.+
Both are giant ionic lattices, but Mg2+ and O2- carry double charges and are smaller ions, so the electrostatic attraction between ions is much stronger and more energy is needed to break the lattice.
3. Predict the shape and bond angle of BF3 and of NH3, and explain the difference.+
BF3 has three bonding pairs and no lone pairs, so it is trigonal planar at 120 degrees. NH3 has three bonding pairs and one lone pair; the lone pair repels more strongly, giving a trigonal pyramidal shape at about 107 degrees.
4. Describe how a dative covalent bond forms in the ammonium ion.+
The lone pair on the nitrogen atom of ammonia is donated to an H+ ion, which has no electrons. Both shared electrons come from nitrogen, and all four bonds are then identical.
5. Explain why ammonia boils at a much higher temperature than phosphine, PH3.+
Ammonia molecules hydrogen bond, because H is bonded to a small, highly electronegative N with a lone pair. Phosphine has only permanent dipole-dipole and London forces, which are weaker, so less energy is needed to separate the molecules.
6. State whether each of CO2, H2O and CHCl3 has an overall dipole, and justify each answer.+
CO2 is linear so the two polar bonds cancel and it has no overall dipole. H2O is bent, so its dipoles do not cancel and it is polar. CHCl3 is tetrahedral but not symmetrical in its atoms, so the dipoles do not cancel and it is polar.
7. Explain why graphite is used as a lubricant and as an electrode.+
Weak forces between layers let them slide over one another, and one delocalised electron per carbon atom can move along the layers to carry charge.
8. Explain in terms of intermolecular forces why boiling points increase from F2 to I2.+
The number of electrons increases down the group, so the induced dipoles and therefore the London forces between molecules get stronger and require more energy to overcome.
Chemical Bonding flashcards and quizzes
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Example flashcards
Q: Why do ionic compounds conduct electricity when molten but not as solids?
A: When molten the ions are free to move and carry charge. In the solid lattice the ions are held in fixed positions.
Q: What type of bonding is present in diamond, and why is it so hard?
A: Giant covalent, each carbon is bonded to four others by strong covalent bonds in a rigid 3D lattice.
Q: Define electronegativity.
A: The ability of an atom to attract the pair of electrons in a covalent bond towards itself.
Q: Why does water have a higher boiling point than expected for its molecular mass?
A: Water molecules form hydrogen bonds between the lone pair on O and the H of another molecule, these are much stronger than the dipole-dipole or London forces of similar-sized molecules.
Quick summary
Chemical Bonding is a high-yield Chemistry topic for GCSE and A-Level students (AQA, Edexcel, OCR). At A-Level you add electronegativity, polarity and the three intermolecular forces (London, permanent dipole-dipole, hydrogen bonding). Examiners reward precise definitions and applied explanations, focus on the core ideas and the small set of terms that come up every series.
Key terms
- Ionic bond
- Covalent bond
- Metallic bond
- Electronegativity
- Polar bond
- Hydrogen bonding
- London forces
- Lattice
- VSEPR
Chemical Bonding FAQs
Why does graphite conduct electricity but diamond does not?+
In graphite each carbon forms three bonds, leaving one delocalised electron per atom that can move along the layers. In diamond every electron is in a localised bond, so none are free to carry charge.
What is the difference between a polar bond and a polar molecule?+
A polar bond has a permanent dipole due to a difference in electronegativity. A polar molecule has a net dipole, symmetrical molecules (like CO2) can have polar bonds but no overall dipole.
Why are noble gases monatomic?+
They have full outer shells, so they have no tendency to gain, lose or share electrons. They exist as single atoms with only weak London forces between them.
Why are metals malleable?+
The layers of positive ions can slide over each other without breaking the metallic bond, because the sea of delocalised electrons holds them together regardless of position.
What is the difference between intramolecular bonds and intermolecular forces?+
Intramolecular bonds, such as covalent bonds, hold atoms together inside a molecule. Intermolecular forces act between separate molecules and are much weaker, so they are what you overcome when a simple molecular substance melts or boils.
How much bonding do I need for A-Level Chemistry compared with GCSE?+
The GCSE ideas of ionic, covalent and metallic bonding and giant versus simple structures still apply. A-Level adds dative covalent bonding, electronegativity and polarity, electron-pair repulsion shapes and bond angles, and the three intermolecular forces, and expects far more precise explanation language.
What is the quickest way to revise chemical bonding?+
Work through one concept at a time, then test yourself out loud on definitions and explanations rather than rereading. Turning your notes into flashcards or a quiz and revisiting them over several days is far more effective than a single long session.
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