VCE Chemistry
Lewis Structures and Covalent Bonding: A VCE Chemistry Guide
Lewis structures trip up more VCE students than almost any other Unit 1 topic, not because the idea is complicated, but because small errors (a missing lone pair, a forgotten double bond) cost marks that are genuinely easy to earn. VCAA examiners award marks for precision, so knowing exactly how to draw and interpret these structures for the ten prescribed molecules is one of the most reliable ways to bank marks on the VCAA Chemistry Examination.
What the syllabus asks
For VCE Unit 1 Chemistry, you need to be able to draw and interpret Lewis (electron dot) structures, structural formulas, and molecular formulas for a specific set of molecules: hydrogen (H2), oxygen (O2), chlorine (Cl2), nitrogen (N2), hydrogen chloride (HCl), carbon dioxide (CO2), water (H2O), ammonia (NH3), methane (CH4), ethane (C2H6), and ethene (C2H4). You also need to understand what covalent bonding is and how to identify bonding pairs and lone pairs in a finished structure.
The idea, explained
What is a covalent bond?
A covalent bond forms when two non-metal atoms share one or more pairs of valence electrons. The bond itself is the electrostatic attraction between that shared electron pair and the positive nuclei of both atoms. Sharing can involve one pair (a single bond), two pairs (a double bond), or three pairs (a triple bond).
When the two atoms sharing electrons have similar electronegativities, the sharing is roughly equal and the bond is non-polar covalent. When there is a moderate difference in electronegativity, one atom pulls the electrons closer to itself, creating a polar covalent bond with a slightly negative end (delta minus) and a slightly positive end (delta plus). HCl is a good example: the electronegativity difference is approximately 1.0, so the Cl end carries a delta minus charge and the H end carries a delta plus charge.
What does a Lewis structure show?
A Lewis structure displays every valence electron in a molecule. Shared (bonding) pairs are drawn as lines between atoms; non-bonding (lone) pairs sit as two dots on a single atom. The guiding rule is that each hydrogen atom needs two electrons (a duet) and each second-row atom such as carbon, nitrogen, oxygen, or chlorine needs eight electrons (an octet).
How to draw one, step by step
- Count all valence electrons by adding the group numbers for every atom.
- Identify the central atom. Hydrogen is never central; choose the least electronegative non-hydrogen atom.
- Draw single bonds from the central atom to each surrounding atom (each bond uses two electrons).
- Spread remaining electrons as lone pairs on outer atoms first, then on the central atom, aiming for octets on non-hydrogen atoms.
- If the central atom still lacks an octet, convert lone pairs on neighbouring atoms into double or triple bonds.
The ten prescribed molecules
Here is what each molecule looks like and what to count:
- H2: H-H. One single bond, no lone pairs. Each H has its duet.
- O2: O=O. One double bond, two lone pairs on each O. Each O has its octet.
- Cl2: Cl-Cl. One single bond, three lone pairs on each Cl. Each Cl has its octet.
- N2: N triple bond N. One triple bond, one lone pair on each N. Each N has its octet.
- HCl: H-Cl. One single bond, three lone pairs on Cl. H has its duet, Cl has its octet. Polar covalent bond.
- CO2: O=C=O. Two C=O double bonds, two lone pairs on each O, no lone pairs on C. C has its octet.
- H2O: H-O-H. Two single bonds, two lone pairs on O. O has its octet, each H has its duet.
- NH3: Three N-H single bonds, one lone pair on N. N has its octet, each H has its duet.
- CH4: Four C-H single bonds, no lone pairs on C. C has its octet, each H has its duet.
- C2H6 (ethane): One C-C single bond and six C-H single bonds. Each C bonds to three H atoms and one C atom; no lone pairs on either C. Each C has its octet.
- C2H4 (ethene): One C=C double bond and four C-H single bonds. Each C bonds to two H atoms and the other C via a double bond; no lone pairs on either C. Each C has its octet.
The key contrast between ethane and ethene is that double bond. In ethene, each carbon uses a double bond to the other carbon to reach its octet with only two hydrogen atoms attached, whereas in ethane a single bond between the carbons is enough because each carbon also bonds to three hydrogen atoms.
What the exam asks
VCAA examiners commonly ask you to draw the Lewis structure of one of these molecules, label bonding pairs and lone pairs, identify the type of bond (single, double, or triple), or explain whether a bond is polar or non-polar covalent. You may also be asked to write the molecular formula or structural formula alongside the Lewis structure. Practise drawing all ten molecules from memory, counting pairs carefully each time.
Common mistakes
- Forgetting lone pairs on outer atoms. In HCl, for example, students often draw only the H-Cl bond and omit the three lone pairs on chlorine. VCAA examiners expect every valence electron to be accounted for.
- Leaving the central atom short of an octet without converting to a double or triple bond. In CO2, a student who draws only single bonds (O-C-O) leaves carbon with only four electrons. The lone pairs on each oxygen must be converted into a second bond to give C its octet.
- Confusing bonding pairs and lone pairs. A double bond contains two bonding pairs, not one. Count each line as one bonding pair.
- Placing hydrogen in the centre. Hydrogen can only ever form one bond, so it always sits on the outside of a structure.
- Mixing up ethane and ethene. Ethane (C2H6) has a C-C single bond; ethene (C2H4) has a C=C double bond. The molecular formulas are your first clue, so check the hydrogen count before you draw anything.
If you want to work through every one of these molecules with immediate feedback and step-by-step guidance, Avocado's AI-native tutoring service has you covered. Start with the Lewis Structures and Covalent Bonding lessons and build the confidence to draw these structures correctly every time under exam conditions.