Simple Molecules and Covalent BondsCambridge IGCSE Chemistry: Revision notes
Section 1
What is a covalent bond and how does it form?
A covalent bond is formed when a pair of electrons is shared between two atoms. This sharing of electrons allows both atoms to achieve noble gas electronic configurations, making them more stable.
Covalent bonding typically occurs between non-metal atoms. Each atom contributes one electron to the shared pair, and both atoms benefit from the extra electrons in their outer shell.
Key principle: Atoms form covalent bonds to gain the stable electron arrangement of noble gases (full outer electron shells).
Examiners always expect you to explain that atoms form covalent bonds to achieve noble gas configurations. Simply saying 'electrons are shared' is not sufficient — you must explain the purpose.
Section 2
How do we represent covalent bond formation using dot-and-cross diagrams?
Dot-and-cross diagrams (also called Lewis structures) show the electrons in the outer shells of atoms before and after bonding.
How to construct them:
- Draw the outer shell electrons of each atom (use dots for one atom and crosses for the other)
- Identify electrons that are unpaired (available for bonding)
- Show the shared pairs in the overlapping region between atoms
- Ensure both atoms now have a stable outer shell
Single covalent bonds (one shared pair):
- H₂: Each hydrogen atom shares one electron; both atoms achieve the helium configuration (2 electrons)
- Cl₂: Each chlorine atom shares one electron; both atoms achieve the argon configuration (8 electrons in outer shell)
- HCl: Hydrogen and chlorine share one electron pair; H achieves helium configuration, Cl achieves argon configuration
Multiple covalent bonds:
- O₂: Oxygen atoms share two electron pairs (double bond); each oxygen achieves the neon configuration
- N₂: Nitrogen atoms share three electron pairs (triple bond); each nitrogen achieves the neon configuration
- C₂H₄: Carbon atoms share two electron pairs (double bond); each carbon also forms single bonds with hydrogen atoms
Molecules with multiple single bonds:
- H₂O: Oxygen forms two single covalent bonds with two hydrogen atoms; oxygen achieves neon configuration
- CH₄: Carbon forms four single covalent bonds with four hydrogen atoms; carbon achieves neon configuration
- NH₃: Nitrogen forms three single covalent bonds with three hydrogen atoms; nitrogen achieves neon configuration
- CH₃OH: Carbon forms four single bonds (three with H, one with O); oxygen forms one bond with C and one with H; all atoms achieve noble gas configurations
- CO₂: Carbon forms two double covalent bonds with two oxygen atoms; all atoms achieve noble gas configurations
For H₂: Each H atom has one electron in its outer shell. When they bond, they share a pair of electrons. Both atoms now have 2 electrons (the helium configuration), so the bond forms. In the dot-and-cross diagram, the shared pair sits between the two H atoms.
For NH₃: Nitrogen has 5 outer electrons and forms three single bonds with three hydrogen atoms. Each N-H bond uses one of nitrogen's electrons and one of hydrogen's electron. After bonding, nitrogen has 8 electrons in its outer shell (achieving neon configuration) and each H has 2 electrons (achieving helium configuration).
Students often forget that both atoms need to achieve a noble gas configuration. If your diagram shows only one atom with a full outer shell, you have made an error. Check every bonded atom.
Section 3
Why do simple molecular compounds have low melting and boiling points?
Simple molecular compounds are composed of small molecules held together by weak intermolecular forces (van der Waals forces or London dispersion forces).
Why the melting and boiling points are low:
- Covalent bonds are strong, but intermolecular forces are very weak
- During melting or boiling, intermolecular forces are broken, not covalent bonds
- Breaking weak intermolecular forces requires very little energy
- Therefore, only small amounts of energy are needed to melt or boil these compounds
- This results in low melting points and low boiling points
Key distinction:
| Property | Covalent Bonds | Intermolecular Forces |
|---|---|---|
| Strength | Very strong | Very weak |
| Location | Within molecules (between atoms) | Between molecules |
| What breaks during melting/boiling | Not broken | Broken |
| Energy required | High | Low |
Examples of low melting/boiling points:
- H₂O: melting point 0 °C, boiling point 100 °C
- CO₂: sublimes at −78 °C (solid directly to gas)
- CH₄: melting point −182 °C, boiling point −162 °C
Think of molecules like groups of friends holding hands (covalent bonds). The hands holding within each group are very strong. But between groups, there are only weak rope connections (intermolecular forces). When heating, the weak ropes snap easily, allowing groups to separate — so little energy is needed.
Examiners want to see that you explain which forces break during melting/boiling. The key phrase is: 'intermolecular forces (not covalent bonds) are broken, which requires little energy.' This distinction is essential for full marks.
Section 4
Why do simple molecular compounds conduct electricity poorly?
Simple molecular compounds are poor electrical conductors because they lack freely moving charged particles.
Why they do not conduct electricity:
- Electrical conduction requires free moving charged particles (ions or free electrons)
- Simple molecular compounds are composed of neutral molecules (not charged ions)
- The electrons are localised (fixed in position within covalent bonds)
- There are no free electrons available to carry electric current
- There are no mobile ions in simple molecular compounds (unlike ionic compounds or molten salts)
- Therefore, electric current cannot flow through simple molecular compounds
Exception: Some simple molecular compounds can conduct electricity when dissolved in water or melted, because they may ionise or the solvent provides mobile ions. However, the compounds themselves in solid or pure liquid form do not conduct.
| State | Conducts Electricity? | Reason |
|---|---|---|
| Solid simple molecular compound | No | No mobile ions or free electrons |
| Molten simple molecular compound | Generally no | Still no ions (unless it ionises) |
| Dissolved in water | Possibly yes | Only if the compound ionises in solution |
| Aqueous ionic solution | Yes | Ions are mobile in solution |
When explaining poor conductivity, always mention that electrons are localised in covalent bonds and there are no free/mobile ions. Avoid vague statements like 'molecules can't conduct' — examiners want evidence of understanding the structure.
Section 5
What are the key structural features that determine properties of simple molecules?
The properties of simple molecular compounds are directly linked to their structure:
Structural features:
- Covalent bonds within molecules: Strong and determine molecular shape and stability
- Intermolecular forces between molecules: Weak and determine bulk properties (melting point, boiling point, solubility)
- Electron localisation: Electrons are fixed within covalent bonds; no free electrons or mobile ions
- Molecular size: Larger molecules often have stronger intermolecular forces and higher melting/boiling points
How structure determines properties:
| Property | Determined by | Reason |
|---|---|---|
| Melting/Boiling point | Intermolecular forces | Weak forces need little energy to break |
| Electrical conductivity | Electron localisation and lack of ions | No mobile charge carriers |
| Solubility in water | Polarity and hydrogen bonding | Varies with molecular structure |
| State at room temperature | Intermolecular force strength | Weak forces → often gases or liquids |
Comparison of common simple molecules:
| Molecule | Bonding | Intermolecular Forces | Melting Point (°C) | Boiling Point (°C) |
|---|---|---|---|---|
| H₂ | Single H-H bond | Very weak | −259 | −253 |
| H₂O | Single O-H bonds | Hydrogen bonds (strong) | 0 | 100 |
| CO₂ | Double C=O bonds | Very weak | −78 (sublimes) | — |
| CH₄ | Single C-H bonds | Very weak | −182 | −162 |
| N₂ | Triple N≡N bond | Very weak | −210 | −196 |
Must Know
- A covalent bond is formed when atoms share a pair of electrons to achieve noble gas electronic configurations (full outer shells)
- Dot-and-cross diagrams show outer shell electrons before and after bonding; use dots for one atom and crosses for the other
- Common molecules with single bonds: H₂, Cl₂, H₂O, CH₄, NH₃, HCl, CH₃OH; molecules with multiple bonds: O₂, N₂, C₂H₄, CO₂
- Intermolecular forces (not covalent bonds) are broken during melting and boiling, requiring very little energy — this explains the low melting and boiling points of simple molecular compounds
- Simple molecular compounds are poor electrical conductors because electrons are localised in covalent bonds and there are no free electrons or mobile ions to carry current
- Covalent bonds are strong but intermolecular forces are weak — this is the key principle explaining why simple molecules have low melting points and poor conductivity
That's the notes covered.
Carry on to the next subtopic.