← Unit 1: How can the diversity of materials be explained?
How can the versatility of non-metals be explained?
the nature of covalent bonding, the construction of Lewis (electron-dot) structures, and the use of valence shell electron pair repulsion (VSEPR) theory to predict the shapes and polarity of simple molecules
A focused VCE Chemistry Unit 1 answer on covalent bonding. Covers the formation of covalent bonds, Lewis (electron-dot) structures including for ions, VSEPR-based shape prediction for the common geometries up to six electron pairs, and how shape plus electronegativity decide overall molecular polarity.
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What this dot point is asking
VCAA wants you to describe the covalent bond as a shared pair of electrons, to draw Lewis structures (including for ions and species with multiple bonds or lone pairs), to predict the shape of small molecules using VSEPR theory, and to combine shape with electronegativity to decide whether a molecule is polar overall.
The answer
Covalent bonding
A covalent bond is a shared pair of electrons between two non-metal atoms. Each atom contributes one electron to the shared pair. The shared pair is attracted to both nuclei, holding the atoms together.
Bonds can be single (one shared pair), double (two shared pairs) or triple (three shared pairs). The more shared pairs, the shorter and stronger the bond. Most second-period elements obey the octet rule (8 valence electrons), although H (2 electrons) and B (often 6) are common exceptions, and third-period and lower elements may exceed 8 (expanded octet: PCl5, SF6).
Lewis structures
A Lewis (electron-dot) structure shows every valence electron as a pair of dots or as a line (for a bonding pair). Procedure:
- Count the total number of valence electrons (sum the group numbers; add electrons for anion charge, subtract for cation charge).
- Draw a skeleton with the least electronegative central atom (H and F are never central; C is usually central in organic molecules).
- Place a bonding pair between each pair of bonded atoms.
- Distribute the remaining electrons as lone pairs to satisfy octets (start with the outer atoms).
- If electrons are left short, form double or triple bonds by sharing lone pairs from outer atoms.
For an ion, wrap the structure in square brackets and add the charge as a superscript.
Examples:
| Species | Total valence e | Structure (described) |
|---|---|---|
| H2O | 8 | O with 2 H (single bonds) and 2 lone pairs |
| NH3 | 8 | N with 3 H (single bonds) and 1 lone pair |
| CO2 | 16 | O=C=O, with 2 lone pairs on each O |
| N2 | 10 | N triple bond N, with 1 lone pair on each N |
| OH^- | 8 | Wrap [O-H] with 3 lone pairs on O, charge -1 |
| NH4^+ | 8 | N with 4 H, 0 lone pairs, charge +1 |
VSEPR theory
Valence Shell Electron Pair Repulsion theory states that electron pairs around a central atom repel each other and arrange to maximise their separation. Both bonding pairs and lone pairs count. Multiple bonds count as one electron domain.
The basic shapes for 2 to 6 electron pairs:
| Pairs | Bonding | Lone | Shape name | Bond angle |
|---|---|---|---|---|
| 2 | 2 | 0 | Linear | 180 |
| 3 | 3 | 0 | Trigonal planar | 120 |
| 3 | 2 | 1 | Bent (V-shape) | ~118 |
| 4 | 4 | 0 | Tetrahedral | 109.5 |
| 4 | 3 | 1 | Trigonal pyramidal | ~107 |
| 4 | 2 | 2 | Bent (V-shape) | ~104.5 |
| 5 | 5 | 0 | Trigonal bipyramidal | 90 and 120 |
| 6 | 6 | 0 | Octahedral | 90 |
Lone pairs repel more strongly than bonding pairs (they are held by only one nucleus, so they spread out more). The bond angle shrinks slightly each time a bonding pair is replaced by a lone pair (CH4 109.5, NH3 ~107, H2O ~104.5).
Polarity
A bond is polar if the bonded atoms differ in electronegativity. The more electronegative atom carries a partial negative charge and the other a partial positive charge. A common rule of thumb: a difference in Pauling electronegativity of about 0.4 or more produces a noticeably polar bond; about 1.7 or more usually indicates an ionic bond.
A molecule is polar overall if the bond dipoles do not cancel. Two requirements for cancellation:
- The polar bonds are symmetrically arranged (linear, trigonal planar, tetrahedral with identical outer atoms, etc.).
- The outer atoms are all the same.
| Molecule | Shape | Bonds polar? | Symmetric? | Overall polar? |
|---|---|---|---|---|
| CO2 | Linear | Yes | Yes (180 apart) | No |
| H2O | Bent | Yes | No (lone pairs break symmetry) | Yes |
| CH4 | Tetrahedral | Slight | Yes | No |
| CHCl3 | Tetrahedral | Yes | No (different outer atoms) | Yes |
| NH3 | Trigonal pyramidal | Yes | No (lone pair) | Yes |
| BF3 | Trigonal planar | Yes | Yes (120 apart) | No |
| HCl | Linear (2 atoms) | Yes | n/a | Yes |
The shortcut for many VCE questions: if the central atom has lone pairs, the molecule is almost always polar (the lone pair breaks symmetry). If the central atom has no lone pairs and all outer atoms are identical, the molecule is almost always non-polar.
Worked example
Predict the shape and polarity of CH2Cl2 (dichloromethane).
Lewis structure: C in the centre, bonded to 2 H and 2 Cl via single bonds; 0 lone pairs on C.
Electron pairs around C: 4. Shape: tetrahedral.
Are the bond dipoles symmetric? No. C-Cl bonds are strongly polar (Cl is more electronegative than C). C-H bonds are only weakly polar. The two C-Cl dipoles do not cancel because the two C-H bonds are not "equivalent" to them. Net result: polar molecule with the negative end towards the two Cl atoms.
Common traps
Forgetting lone pairs in the VSEPR count. H2O has 4 electron pairs around O, not 2.
Calling CO2 polar. Each C=O is polar, but the two dipoles point opposite ways and exactly cancel. CO2 is non-polar.
Counting a double bond as 2 electron domains. VSEPR counts a multiple bond as a single domain. CO2 has 2 domains around C and is linear.
Drawing a Lewis structure with the wrong number of electrons. Always sum group numbers and adjust for charge before placing electrons.
Forgetting brackets around an ionic Lewis structure. Ions must show the structure in [ ] with the overall charge as a superscript.
Applying octet rules to H and Be/B. Hydrogen wants 2 electrons. Beryllium and boron are commonly stable with 4 and 6, respectively.
In one sentence
Covalent bonds are shared pairs of electrons that you can map with Lewis structures; VSEPR theory predicts the shape from the number of bonding and lone pairs around each central atom; and a molecule is polar overall if its polar bond dipoles do not cancel by symmetry.
Past exam questions, worked
Real questions from past VCAA papers on this dot point, with our answer explainer.
2024 VCE4 marksFor each of NH3, CO2 and H2O: (a) draw the Lewis structure, (b) predict the shape using VSEPR theory, (c) state whether the molecule is polar overall.Show worked answer →
A 4-mark answer needs the right number of bonding/lone pairs, the correct shape name, and a clear polar/non-polar verdict with reasoning.
NH3 (ammonia): N has 3 bonding pairs and 1 lone pair (4 electron pairs total). VSEPR shape with one lone pair: trigonal pyramidal. N is more electronegative than H so each N-H bond is polar; the lone pair breaks the symmetry, the dipoles do not cancel. Polar overall.
CO2 (carbon dioxide): C has 2 double bonds and 0 lone pairs (2 electron domains). VSEPR shape: linear. Each C=O bond is polar (O more electronegative), but the two dipoles point in opposite directions and exactly cancel. Non-polar overall.
H2O (water): O has 2 bonding pairs and 2 lone pairs (4 electron pairs total). VSEPR shape with two lone pairs: bent (or V-shape, ~104.5 degrees). O-H bonds are polar; the lone pairs prevent the dipoles from cancelling. Polar overall.
The recurring trap is calling H2O linear because it has two bonds. Lone pairs count toward the geometry: H2O has 4 electron pairs around O, two are lone, giving a bent shape.
2025 VCE3 marksPredict and justify the shape and bond angle for the BF3 and PCl5 molecules using VSEPR theory.Show worked answer →
A 3-mark answer needs the electron-pair count, the shape name, and the bond angle.
BF3: B has 3 bonding pairs and 0 lone pairs (3 electron pairs total). Electron pairs repel and arrange to maximise separation: trigonal planar, F-B-F angle 120 degrees. BF3 is an electron-deficient species (only 6 electrons around B), which is allowed for B in this course.
PCl5: P has 5 bonding pairs and 0 lone pairs (5 electron pairs total). VSEPR arrangement: trigonal bipyramidal. Three equatorial Cl atoms at 120 degrees, two axial Cl atoms at 90 degrees from the equatorial plane (and 180 degrees from each other). P uses an "expanded octet" because it has access to d orbitals.
Markers reward distinguishing the equatorial and axial positions in PCl5.
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