Covalent vs Ionic Bonding
Complete comparison of ionic and covalent bonding for IGCSE Chemistry 0620: formation, structures, properties, and how to identify each type.
Published by IGCSEChemistry.com.my
Chemistry teaching team: K. S. Tan (15+ years teaching IGCSE Chemistry) and Ms Yash (10+ years teaching IGCSE Chemistry) and Ms Kartini (15+ years teaching IGCSE Chemistry).
Mapped to Cambridge IGCSE Chemistry 0620 (2026–2028). Last updated 2026-08-19.
Bonding questions appear on every Paper 2 and Paper 4 in every exam series. The 0620 mark scheme is ruthlessly specific about vocabulary: “transfer” for ionic, “shared pair” for covalent, “electrostatic attraction” for the force, and “between oppositely charged ions” to complete the ionic bond definition. Swapping any word typically loses the mark.
How each bond forms
Ionic bonding
Ionic bonds form between metals and non-metals. The metal atom loses its outer-shell electrons, forming a positive ion (cation). The non-metal atom gains those electrons, forming a negative ion (anion). The electrostatic attraction between the oppositely charged ions is the ionic bond.
Example: sodium chloride
- Na (2,8,1) loses 1 electron to become Na+ (2,8)
- Cl (2,8,7) gains 1 electron to become Cl- (2,8,8)
- Both ions now have a complete outer shell (the electron configuration of a noble gas)
Covalent bonding
Covalent bonds form between non-metal atoms. Each atom contributes one electron to form a shared pair. By sharing, both atoms achieve a full outer shell.
Example: water (H2O)
- Oxygen (2,6) needs 2 more electrons for a full outer shell
- Each hydrogen (1) needs 1 more electron
- Oxygen shares one pair with each hydrogen, forming two covalent bonds
Dot-and-cross diagrams
These diagrams show only outer-shell electrons, using dots for one atom and crosses for the other. Examiners check:
- Correct number of outer-shell electrons on each atom
- Shared pairs shown in the overlap region
- Lone pairs shown on the correct atom
- Ions shown in square brackets with the charge
Practise drawing diagrams for NaCl, MgO, MgCl2 (ionic) and H2O, CH4, NH3, CO2, HCl, H2, Cl2, O2, N2 (covalent). These cover every variant the exam uses.
Structures: giant vs simple
| Feature | Ionic compounds | Simple covalent molecules | Giant covalent structures |
|---|---|---|---|
| Example | NaCl, MgO | H2O, CO2, CH4 | Diamond, graphite, SiO2 |
| Basic unit | Ions in a lattice | Discrete molecules | Atoms in a network |
| Bonding within | Ionic bonds | Covalent bonds | Covalent bonds |
| Forces between units | Same ionic bonds (continuous lattice) | Weak intermolecular forces | Same covalent bonds (continuous) |
| Melting point | High | Low | Very high |
The distinction between forces within molecules (strong covalent bonds) and forces between molecules (weak intermolecular forces) is the most examined concept in this area. When a simple covalent substance melts, the weak intermolecular forces break. The covalent bonds do not break.
Property comparison
| Property | Ionic | Simple covalent | Giant covalent |
|---|---|---|---|
| Melting point | High | Low | Very high |
| Boiling point | High | Low | Very high |
| Conducts as solid? | No (ions fixed) | No (no ions/electrons) | No (except graphite) |
| Conducts when molten? | Yes (ions free to move) | No | No (does not melt easily) |
| Conducts in solution? | Yes (if soluble) | No | Insoluble |
| Soluble in water? | Many are | Some (e.g. sugar) | No |
| Brittle or malleable? | Brittle | Soft / low force to separate | Very hard |
Explaining each property
High melting point (ionic): Strong electrostatic attraction between oppositely charged ions throughout the giant ionic lattice. A large amount of energy is needed to overcome these forces.
Low melting point (simple covalent): The intermolecular forces between molecules are weak. Little energy is needed to overcome them. The strong covalent bonds within molecules do not break during melting.
Conduction (ionic): In the solid, ions are in fixed positions and cannot move. When molten or dissolved, ions are free to move and carry charge. See why ionic compounds conduct.
No conduction (simple covalent): No ions or free electrons exist in the substance. There are no mobile charge carriers.
Graphite conducts: Each carbon bonds to three others; the fourth outer electron is delocalised. These delocalised electrons carry charge, similar to metallic bonding.
How to identify bonding type from properties
A common Paper 2 question gives you properties and asks you to identify the bonding type:
- High melting point + conducts when molten = ionic
- High melting point + does not conduct = giant covalent
- Low melting point + does not conduct = simple covalent
- High melting point + conducts as solid = metallic
This is a decision tree examiners love. Learn it.
Intermediate cases examined in 0620
Hydrogen chloride (HCl): A covalent molecule. Pure HCl gas does not conduct. But dissolved in water it ionises completely (HCl -> H+ + Cl-), producing ions that conduct. The compound is covalent, but its aqueous solution behaves like an ionic solution.
Ammonium compounds (e.g. NH4Cl): The ammonium ion NH4+ contains covalent bonds (N-H), but ammonium chloride is an ionic compound (NH4+ and Cl- ions). It has ionic properties: high melting point, conducts when dissolved.
Common exam mistakes
- “Ionic bonds are stronger than covalent bonds” — not necessarily true and not what the syllabus says. Both are strong. The key difference is what breaks on melting (intermolecular forces for simple covalent vs ionic bonds for ionic compounds).
- “Covalent compounds have low melting points” — only true for simple covalent molecules. Giant covalent structures (diamond, SiO2) have very high melting points.
- “NaCl has molecules” — ionic compounds do not form molecules. They form a giant ionic lattice. The formula NaCl gives the simplest ratio of ions.
- Using “sharing” for ionic bonds — ionic is transfer, covalent is sharing. These verbs are non-negotiable in mark schemes.
Worked exam questions
The table below shows properties of two substances, A and B.
Melting point Conducts as solid Conducts when molten A High No Yes B Low No No
Identify the type of bonding and structure in A and B. Explain your answers. [4 marks]
- A has ionic bonding and a giant ionic structure [1]. High melting point due to strong electrostatic attraction between ions; does not conduct as solid because ions are in fixed positions; conducts when molten because ions are free to move [1]
- B has covalent bonding and a simple molecular structure [1]. Low melting point because weak intermolecular forces need little energy to overcome; no conduction because no ions or free electrons [1]
Draw a dot-and-cross diagram for magnesium oxide, MgO. Show the charges on the ions. [3 marks]
- Mg shown losing 2 outer electrons (2,8,2 -> 2,8), Mg2+ ion in square brackets with 2+ charge [1]
- O shown gaining 2 electrons (2,6 -> 2,8), O2- ion in square brackets with 2- charge [1]
- Electrons transferred shown correctly (dots on Mg become crosses on O or vice versa, showing which came from where) [1]
Explain why sodium chloride has a high melting point but methane has a low melting point. [4 marks]
- NaCl has a giant ionic structure with strong electrostatic attraction between Na+ and Cl- ions [1]
- A large amount of energy is needed to overcome these strong ionic bonds [1]
- CH4 is a simple covalent molecule with weak intermolecular forces between molecules [1]
- Little energy is needed to overcome these weak forces (the strong C-H covalent bonds within molecules do not break) [1]
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