Structure and Bonding Revision Guide

Structure and bonding is one of the most important topics in chemistry because it explains properties: why salt melts at 801 °C, why water is a liquid, why copper conducts and why diamond is hard. It’s also one of the most examined. This revision guide gathers the essentials into checklists, summary tables and exam tips. Work through each section, ticking off what you can explain without notes, and follow the links for anything that needs more work.

1. The three types of bonding

• Define a chemical bond as an electrostatic attraction that holds atoms or ions together (see what is a chemical bond?).
• State which elements form each type of bond.
• Explain what happens to electrons in each type.

Bonding	Between	Electrons	Particles
Ionic	Metal + non-metal	Transferred	Positive and negative ions
Covalent	Non-metal + non-metal	Shared in pairs	Atoms in molecules or networks
Metallic	Metal atoms	Delocalised “sea”	Positive ions + delocalised electrons

Key definitions to learn word-for-word:

• Ionic bond: the electrostatic attraction between oppositely charged ions.
• Covalent bond: a shared pair of electrons between two atoms.
• Metallic bond: the electrostatic attraction between positive metal ions and delocalised electrons.

2. Ionic bonding checklist

• Work out ion charges from the group number: groups 1, 2, 3 → +1, +2, +3; groups 5, 6, 7 → −3, −2, −1 (see ionic bonding explained).
• Write formulas by balancing charges: MgO, CaCl₂, Al₂O₃, Na₂O.
• Draw dot-and-cross diagrams for ionic compounds with square brackets and charges (see how to draw dot-and-cross diagrams).
• Describe the giant ionic lattice (NaCl: each ion surrounded by 6 of the opposite charge) (see ionic lattices).
• Explain why higher charges and smaller ions give higher melting points.

3. Covalent bonding checklist

• Know the usual number of bonds: H 1, O 2, N 3, C 4, halogens 1 (see covalent bonding explained).
• Draw dot-and-cross diagrams for H₂, Cl₂, HCl, H₂O, NH₃, CH₄, O₂, N₂, CO₂.
• Identify single, double and triple bonds and lone pairs (see single, double and triple bonds compared).
• Explain the difference between simple molecular and giant covalent structures (see simple molecular vs giant covalent).

4. Metallic bonding checklist

• Describe the “positive ions in a sea of delocalised electrons” model (see metallic bonding).
• Explain conductivity, malleability and high melting points (see how metallic bonding explains conductivity, lustre and malleability).
• Explain why alloys are harder than pure metals (see alloys).

5. The four structure types

This is the big table to learn. Almost every structure-and-bonding exam question can be answered from it.

Structure	Examples	Particles	Forces overcome on melting	Melting point	Conducts solid?	Conducts liquid?	Soluble in water?
Giant ionic	NaCl, MgO, CaCO₃	Ions	Strong ionic bonds	High	No	Yes	Often
Simple molecular	H₂O, CO₂, I₂, CH₄	Molecules	Weak intermolecular forces	Low	No	No	Only if polar
Giant covalent	Diamond, graphite, SiO₂	Atoms	Strong covalent bonds	Very high	No (graphite yes)	—	No
Giant metallic	Cu, Fe, Na, alloys	Positive ions + electrons	Metallic bonds	Usually high	Yes	Yes	No (some react)

Using data to identify the structure

Clue	Structure
Low melting point, no conduction	Simple molecular
Very high melting point, no conduction in any state	Giant covalent
High melting point, conducts only when molten or dissolved	Giant ionic
Conducts as solid and liquid, malleable	Metallic
Very high melting point and conducts, but a non-metal	Graphite

6. Properties explained — model sentences

Use these as templates for “explain” questions:

• NaCl has a high melting point because it has a giant ionic lattice with strong electrostatic attractions between oppositely charged ions in all directions, which need a lot of energy to overcome.
• Water has a low boiling point (compared with ionic compounds) because it’s a simple molecular substance; only the weak intermolecular forces (hydrogen bonds) between molecules are overcome, not the covalent bonds.
• Diamond is very hard because each carbon is covalently bonded to four others in a rigid giant network.
• Graphite conducts electricity because each carbon forms only three bonds, leaving one delocalised electron per atom that can move along the layers (see giant covalent structures).
• Copper conducts electricity because it has delocalised electrons that can move through the structure.
• Metals are malleable because layers of positive ions can slide over each other while the delocalised electrons keep holding them together.
• Ionic compounds are brittle because a force shifts layers so that ions of the same charge line up and repel.
• Solid NaCl doesn’t conduct because the ions are fixed in the lattice; molten or dissolved NaCl does because the ions can move.

7. Intermolecular forces (for simple molecules)

• Distinguish bonds (within molecules) from intermolecular forces (between molecules) (see intermolecular forces).
• Know the three main types: London (dispersion) forces (all molecules; stronger for bigger molecules), permanent dipole–dipole (polar molecules) and hydrogen bonds (H bonded to N, O or F).
• Explain boiling-point trends, e.g. halogens rise down the group; water is anomalously high (see hydrogen bonding).

8. Shapes and polarity (higher level)

• Use electron pairs to predict shapes: 2 → linear 180°; 3 → trigonal planar 120°; 4 → tetrahedral 109.5° (see VSEPR and molecular geometry).
• Lone pairs: NH₃ pyramidal 107°, H₂O bent 104.5° (see methane, ammonia and water compared).
• Decide whether a molecule is polar from bond polarity and shape (see how to tell if a molecule is polar).

9. Ten quick-fire questions

1. What type of bonding is in potassium bromide?
2. Formula of the compound formed from Al³⁺ and O²⁻?
3. Why does CO₂ have a low boiling point?
4. Name a giant covalent substance that conducts electricity.
5. What moves when molten NaCl conducts?
6. What moves when copper conducts?
7. How many covalent bonds does a nitrogen atom usually form?
8. Why are alloys harder than pure metals?
9. What is the shape of methane?
10. Is solid magnesium oxide a conductor?

Answers: 1 ionic; 2 Al₂O₃; 3 simple molecular — weak intermolecular forces between molecules; 4 graphite; 5 ions; 6 delocalised electrons; 7 three; 8 different-sized atoms distort the layers so they can’t slide easily; 9 tetrahedral; 10 no — ions can’t move in the solid.

10. Exam technique

• Name the structure first, then the particles, then the forces, then link to the property.
• Never say “molecules” for ionic compounds or metals.
• Never say “breaking covalent bonds” when a simple molecular substance melts or boils.
• Say “delocalised electrons”, not “free electrons” alone — many mark schemes want “delocalised”.
• Say “strong electrostatic attraction” for ionic and metallic bonding — and “a lot of energy is needed to overcome”.
• Compare like with like: when asked why MgO melts higher than NaCl, mention higher charges on the ions.

Key takeaways

• Ionic (transfer), covalent (sharing) and metallic (delocalised electrons) bonding.
• Four structures: giant ionic, simple molecular, giant covalent, giant metallic.
• What is overcome on melting decides the melting point: weak intermolecular forces (low) vs strong bonds (high).
• Conductivity needs mobile charged particles: ions (molten/dissolved ionic) or delocalised electrons (metals, graphite).

Check your understanding with chemical bonding misconceptions.