3i. Covalent bonding in the methane molecule CH₄

Covalent bonding diagram for METHANE covalent molecule, molecular formula CH₄

Four atoms of hydrogen (1) combine with one atom of carbon (2.4) to form the molecule of the compound methane CH₄

Each hydrogen atom is one electron short of a helium structure (full shell) and carbon is four electrons short of a full outer shell (of 8), so four hydrogen atoms share their electrons with the four outer electrons of carbon, so all five atoms effectively have full outer shells.

four of  and one of combine to form so that the hydrogen atoms are electronically like helium and the carbon atom becomes like neon (only the outer shell of carbon's bonding electrons are shown).

Electronically, hydrogen (1) becomes like helium (2) and carbon (2.4) becomes like neon (2.8), so the hydrogen and carbon atoms effectively have full outer shells in forming the covalent bonds when the atoms share their outer electrons.

(Lewis diagram of methane) simplified 'dot and cross' electronic diagram for the covalently bonded methane molecule

The methane molecule is held together by the four strong C–H carbon–hydrogen covalent bonds by sharing electrons.

Note that the inner shell of carbon's electrons are not shown above, only the outer shell of carbon's electrons are involved in the covalent bonding.

The molecule can be shown as  (displayed formula) with four carbon – hydrogen single covalent bonds (A level note: its called a tetrahedral shape, the H–C–H bond angle is 109^o).

SiH₄ will be similar because silicon (2.8.4) is in the same group as carbon. This displayed formula does indicate the shape of the methane molecule as well as how the four single C-H covalent bonds are arranged, but no relative size of atoms or electronic detail of bond formation by electron sharing.

All the bonds in the above examples are single covalent bonds. Below are three examples 7–9, where there is a double bond in the molecule, in order that the atoms have stable Noble Gas outer electron arrangements around each atom. Carbon and silicon have a valency of 4.

More complex examples can be worked out e.g. involving C, H and O. In each case link in the atoms so that there are 2 around a H (electronically like He), or 8 around the C or O (electronically like Ne).

 On the left are full 'dot and cross' electronic Lewis diagram for the covalent bonding in the methane molecule.

 The electronic dot & cross Lewis diagram give a 2D view of the molecule in terms of electrons, but little idea on the shape of the molecule which might be anything but flat!

The Venn diagram style (above left) is the best style for methane, clearly showing the sharing of the pairs of electrons for the single covalent bonds in the chlorine molecule (in a sort of Venn diagram style).

Below, the stick and ball diagram of methane gives a much better impression of the shape of the molecule and the spatial arrangement of the four single bonds and five atoms.

   AND 

ball and stick model of methane  AND    space-filling model of methane

Comments

Melting point of methane -183 ^oC

Boiling point of methane -162 ^oC

You would expect low values because of the weak intermolecular forces between small covalent molecules like methane.

Methane is a colourless flammable gas at room temperature, it has a strong 'hydrocarbon odour'.

An important fossil fuel.

Extra notes about Methane (CH₄) e.g. for GCSE/IGCSE Level Chemistry Revision Notes

Properties of methane

• Formula: CH₄
• Structure: Tetrahedral, covalent molecule (bond angle ≈ 109.5°)
• Physical properties:
  • Colourless, odourless gas (odour added for domestic use)
  • Less dense than air
  • Insoluble in water
• Chemical properties:
  • Burns with a pale blue flame → carbon dioxide + water
    CH₄ + 2O₂ → CO₂ + 2H₂O
  • Highly flammable
  • Greenhouse gas (contributes to climate change)

Reactions of methane

• Complete combustion: CH₄ + 2O₂ → CO₂ + 2H₂O (blue flame, maximum energy release)
• Incomplete combustion: Produces carbon monoxide (CO) or carbon (soot)
  • CH₄ + O₂ → CO + 2H₂O
  • CH₄ + O₂ → C + 2H₂O
• Substitution reactions: With halogens in UV light
  • CH₄ + Cl₂ → CH₃Cl + HCl (chloromethane)
• Steam reforming: CH₂ + H₂O → CO + 3H₂ (important industrial process to make hydrogen)

Industrial Manufacture of methane

• Natural gas extraction: Methane is the main component of natural gas (from fossil fuel deposits).
• Biogas production: Anaerobic decomposition of organic matter (e.g., animal waste, plant material).
• Synthetic processes: Methane can be produced during coal gasification or as a by-product in chemical industries.

Uses of Methane

• Fuel: Domestic heating, cooking, electricity generation.
• Feedstock: Steam reforming → hydrogen (used in Haber process, fuel cells).
• Chemical industry: Production of methanol, chloromethanes, and other organic compounds.
• Environmental context: Methane capture from waste management and agriculture.

Typical Exam Board Requirements

See also

Pollution, carbon monoxide, nitrogen oxides, climate change-global warming (3 relevant pages)

Student Exam Tips

• Equations: Always balance combustion equations correctly.
• Complete versus incomplete combustion: State products clearly (CO₂ versus CO/soot).
• Substitution reaction: Remember UV light is required for halogen substitution.
• Environmental impact: Link methane to greenhouse effect and climate change.
• Industrial context: Steam reforming → hydrogen; biogas → renewable energy.
• Exam wording: Methane is the main component of natural gas — don’t confuse with “natural gas = methane only.”

Common Misconceptions

•  Thinking methane is soluble in water (it is insoluble).
•  Forgetting methane is tetrahedral (students sometimes assume linear).
•  Confusing complete combustion (CO₂ + H₂O) with incomplete combustion (CO or C formed).
•  Assuming methane is odourless in domestic supply (odorants are added for safety).
•  Believing methane is only harmful as a fuel — it is also a potent greenhouse gas.