Covalent Bonding
Summary: Covalent bonding involves sharing of electron pairs between non-metal atoms to achieve a full outer shell (noble gas configuration). Can form simple molecular structures (weak intermolecular forces, low mp/bp) or giant covalent networks (diamond, graphite, SiO₂) with very high melting points. Tags: igcse chemistry bonding covalent structures Created: 2026-07-14 Last Updated: 2026-07-18
Covalent bonding is the type of chemical bonding in which pairs of electrons are shared between non-metal atoms, with each shared pair constituting a single Covalent Bond. By sharing electrons rather than transferring them, each atom gains access to enough outer-shell electrons to attain a Noble Gas Electronic Configuration — two electrons for hydrogen, eight for most other atoms. This bonding produces Simple Molecules such as Hydrogen (H2), Chlorine (Cl2), Oxygen (O2) and Nitrogen (N2), as well as Simple Molecular Compounds like Water (H2O), Methane (CH4), Ammonia, Hydrogen Chloride (HCl), Carbon Dioxide (CO2), Methanol (CH3OH) and Ethene (C2H4). The electron sharing in each of these species can be represented using a Dot-and-Cross Diagram, which shows outer-shell electrons only and highlights shared pairs in the overlap between atoms. While the covalent bonds within molecules are strong, each molecule is attracted to its neighbours only by a weak Intermolecular Force, which explains the characteristically low melting and boiling points of simple molecular substances. In contrast, giant covalent structures such as diamond, graphite and silicon dioxide consist of continuous lattices of covalently bonded atoms, giving them very high melting points.
What is Covalent Bonding?
Covalent bonding occurs between non-metal atoms. Instead of transferring electrons (as in ionic bonding), atoms share pairs of electrons so that each atom achieves a full outer shell — the stable electronic structure of a noble gas (Group 0).
Each shared pair of electrons = one covalent bond. By sharing, each atom can count the shared electrons as part of its own outer shell without permanently losing or gaining electrons.
Example — Hydrogen (H₂): Each H atom has 1 outer electron and needs 1 more (helium configuration, 2 electrons). Two H atoms each contribute 1 electron to form a shared pair. Both now have access to 2 electrons → single covalent bond: H—H.
Key rule: For elements in Groups 4–7, the number of covalent bonds an atom forms = 8 − group number:
- Group 7 (e.g., Cl): 1 bond
- Group 6 (e.g., O): 2 bonds
- Group 5 (e.g., N): 3 bonds
- Group 4 (e.g., C): 4 bonds
Hydrogen is the exception — it follows the “duet rule” (needs 2 electrons, so forms 1 bond).
Single, Double, and Triple Covalent Bonds
| Bond Type | Shared Pairs | Electrons Shared | Example | Notation |
|---|---|---|---|---|
| Single | 1 | 2 | Cl₂, CH₄, H₂ | Cl—Cl |
| Double | 2 | 4 | O₂, CO₂ | O=O |
| Triple | 3 | 6 | N₂ | N≡N |
- Chlorine (Cl₂): Each Cl has 7 outer electrons. Sharing 1 pair → each Cl effectively has 8.
- Oxygen (O₂): Each O has 6 outer electrons. Sharing 2 pairs (double bond) → each O effectively has 8. Written as O=O.
- Nitrogen (N₂): Each N has 5 outer electrons. Sharing 3 pairs (triple bond) → each N effectively has 8. Written as N≡N. This very strong bond explains why N₂ is so unreactive.
- Carbon dioxide (CO₂): C forms two double bonds, one to each O: O=C=O.
Bond strength: triple > double > single — more shared pairs = stronger bond = more energy to break.
Simple Molecular Structures
Substances like H₂, Cl₂, H₂O, CH₄, CO₂, O₂, N₂ exist as simple molecules — small groups of atoms held together by covalent bonds. Within each molecule, atoms are held by strong covalent bonds. Between molecules, only weak intermolecular forces (van der Waals’ forces) act. Most are gases or low-boiling liquids at room temperature.
Properties: low mp/bp (only weak intermolecular forces overcome, covalent bonds stay intact); do not conduct electricity (no mobile charged particles); often insoluble in water (though HCl, NH₃ dissolve and react).
Crucial distinction: When simple molecular substances melt/boil, intermolecular forces are overcome — covalent bonds stay intact.
Giant Covalent Structures
In giant covalent structures, billions of atoms are joined by covalent bonds in a continuous lattice. There are no individual molecules — the entire structure is one giant molecule.
Diamond
- Each carbon atom covalently bonded to 4 other C atoms in a tetrahedral arrangement (bond angle 109.5°)
- All 4 outer electrons per carbon used in C—C single bonds
- Forms a rigid 3D giant covalent lattice
| Property | Explanation |
|---|---|
| Extremely hard | 4 strong covalent bonds per C in rigid 3D lattice — all bonds must be broken to deform |
| Very high mp (~3550°C) | Strong covalent bonds throughout require enormous energy to break |
| Does not conduct | No delocalised electrons — all outer electrons used in bonds |
| Insoluble | Covalent bonds cannot be broken by water |
Uses: Cutting tools, drill bits, jewellery.
Graphite
- Each carbon atom covalently bonded to 3 other C atoms, forming flat hexagonal layers
- The 4th outer electron on each C is delocalised — free to move between layers
- Layers held together by weak intermolecular forces
| Property | Explanation |
|---|---|
| Soft and slippery | Layers held by weak intermolecular forces → slide over each other easily |
| Very high mp (~3650°C) | Strong covalent bonds within each layer require enormous energy to break |
| Conducts electricity | Delocalised electrons (1 per C) carry charge freely |
| Insoluble | Covalent bonds cannot be broken by water |
Uses: Pencil leads, electrodes, lubricant.
Diamond vs graphite comparison: Diamond: each C bonded to 4 others in rigid 3D tetrahedral lattice. Graphite: each C bonded to 3 others in layers; layers held by weak intermolecular forces that allow sliding.
Silicon Dioxide (SiO₂) — Silica
- Each Si atom bonded to 4 O atoms; each O atom bonded to 2 Si atoms
- Continuous 3D tetrahedral lattice — structure similar to diamond
- Formula SiO₂ shows the ratio of atoms (1 Si : 2 O), NOT individual molecules
Properties: Very high mp (~1710°C), hard, does not conduct, insoluble — same reasons as diamond.
Comparison: CO₂ vs SiO₂:
| CO₂ | SiO₂ | |
|---|---|---|
| Structure | Simple molecular | Giant covalent |
| Particles | Small CO₂ molecules | Continuous Si–O lattice |
| Between units | Weak intermolecular forces | Strong covalent bonds (continuous) |
| Melting point | Very low (sublimes at −78°C) | Very high (1710°C) |
| State at RT | Gas | Solid |
Why is CO₂ a gas but SiO₂ a solid? CO₂ is simple molecular — only weak intermolecular forces need to be overcome when boiling, requiring little energy. SiO₂ is giant covalent — many strong covalent bonds must be broken to melt it, requiring enormous energy.
Drawing Dot-and-Cross Diagrams
A dot-and-cross diagram shows outer-shell electrons only, using dots (•) for one atom’s electrons and crosses (×) for the other, with shared pairs in the overlap. Each atom must achieve 8 outer electrons (except H: 2).
Key diagrams for IGCSE:
| Molecule | Bonding Summary |
|---|---|
| H₂ | H—H, 1 shared pair, each H has 2 electrons |
| Cl₂ | Cl—Cl, 1 shared pair, each Cl has 8 outer electrons |
| HCl | H—Cl, 1 shared pair, H has 2, Cl has 8 outer electrons |
| H₂O | O shares 1 pair with each H (2 bonds total), O has 8 (incl. 2 lone pairs), each H has 2 |
| CH₄ | C shares 1 pair with each H (4 bonds total), C has 8, each H has 2 |
| CO₂ | O=C=O, 2 double bonds, each atom has 8 outer electrons |
| O₂ | O=O, 1 double bond, each O has 8 outer electrons |
| N₂ | N≡N, 1 triple bond, each N has 8 outer electrons |
| C₂H₄ | C=C double bond + 2 C—H per carbon, each C has 8, each H has 2 |
| NH₃ | N shares 1 pair with each H (3 bonds), N has 8 (incl. 1 lone pair), each H has 2 |
Intermolecular Forces vs Covalent Bonds
This distinction is fundamental:
| Covalent Bond | Intermolecular Force | |
|---|---|---|
| Location | Within a molecule (between atoms) | Between molecules |
| Strength | Strong — large energy to break | Weak — small energy to overcome |
| On melting/boiling | Do NOT break — molecules stay intact | Are overcome — molecules separate |
| Affects mp/bp? | Only in giant covalent structures | Determines mp/bp of simple molecular substances |
When water boils, only intermolecular forces are overcome. The O—H bonds within each H₂O molecule remain intact.
Key Points
- Covalent bonding = sharing of electron pairs between non-metal atoms to achieve noble gas configuration
- Bonds formed = 8 − group number (Groups 4–7); H forms 1 bond
- Single = 1 shared pair; Double = 2 shared pairs; Triple = 3 shared pairs
- Simple molecular: low mp/bp, no electrical conductivity — weak intermolecular forces between molecules
- Diamond: each C bonded to 4 others, tetrahedral, no delocalised electrons → hard, does not conduct
- Graphite: each C bonded to 3 others, layered, has delocalised electrons → soft/slippery, conducts electricity
- SiO₂: giant covalent, similar structure to diamond, very high mp
- Strong covalent bonds (within molecules) vs weak intermolecular forces (between molecules)
- Melting/boiling simple molecules: intermolecular forces overcome, covalent bonds stay intact
Key Concepts from Past Papers
- Covalent bond: the sharing of a pair of electrons between two atoms
- Diamond: each C bonded to 4 others in a tetrahedral giant covalent lattice
- Graphite: each C bonded to 3 others in layers, with delocalised electrons between layers
- Simple molecular substances have low mp/bp because only weak intermolecular forces need to be overcome
- Giant covalent structures have very high mp because many strong covalent bonds must be broken
Keywords from Past Papers
graphite, reacts, diamond, calcium, oxide, platinum, giant, covalent, down, forms, silicon, silica, high, temperature, catalyst
Related Notes
- Atomic Structure
- Ions and Ionic Bonds
- Metallic Bonding
- Giant Structures
- Organic Chemistry Fundamentals
- Covalent Bond
- Noble Gas Electronic Configuration
- Simple Molecules
- Simple Molecular Compounds
- Dot-and-Cross Diagram
- Intermolecular Force
- Hydrogen (H2)
- Chlorine (Cl2)
- Water (H2O)
- Methane (CH4)
- Ammonia
- Hydrogen Chloride (HCl)
- Methanol (CH3OH)
- Ethene (C2H4)
- Oxygen (O2)
- Carbon Dioxide (CO2)
- Nitrogen (N2)
Sources
- OpenStax Chemistry 2e — Chapter 7: Chemical Bonding and Molecular Geometry, Rice University (free, CC BY 4.0)
- BBC Bitesize GCSE Chemistry — Covalent Bonding, BBC (free educational resource)
- Cambridge IGCSE Chemistry 0620 — Syllabus Section 2.5, Cambridge Assessment International Education
- CK-12 Chemistry for High School — Covalent Bonding, CK-12 Foundation (free, CC BY-NC 3.0)
Past Paper Sources
- 0620/32 Feb/March 2017: Q22(b)(iii) (1m), Q22(d)(i) (1m)
- 0620/32 Feb/March 2020: Q22(b)(ii) (3m)
- 0620/32 Feb/March 2021: Q77(b)(i) (1m)
- 0620/32 Feb/March 2022: Q66(c)(i) (1m)
- 0620/32 May/June 2020: Q55(b)(ii) (1m)
- 0620/33 May/June 2021: Q88(b)(iii) (1m)
- 0620/33 May/June 2024: Q33(a)(ii) (1m)
- 0620/33 October/November 2015: Q44(b)(i) (2m)
- 0620/33 October/November 2017: Q44(f)(i) (0m), Q44(f)(i) (0m)
- 0620/33 October/November 2023: Q55(c)(ii) (1m)
- 0971/31 May/June 2021: Q44(c)(i) (2m)
- 0971/32 May/June 2020: Q55(b)(ii) (1m)
Common Misconceptions
| Misconception | Reality |
|---|---|
| ”Covalent bonds break when water boils” | Only weak intermolecular forces are overcome. O—H bonds within each H₂O remain intact. |
| ”Covalent bonds are weak” | Covalent bonds are strong. The intermolecular forces between molecules are weak. |
| ”Diamond and graphite are different elements” | Both are pure carbon — they are allotropes (different structural forms of the same element). |
| ”Graphite has a low melting point because layers slide” | Graphite has a very high mp (~3650°C) due to strong covalent bonds within layers. Slipperiness is separate. |
| ”SiO₂ has individual molecules like CO₂” | SiO₂ is a giant covalent lattice; the formula only shows the Si:O ratio (1:2). |
| ”All covalent substances have low melting points” | Only simple molecular substances do. Giant covalent structures (diamond, graphite, SiO₂) have very high mp. |
| ”No covalent substance conducts electricity” | Graphite does conduct due to its delocalised electrons. |