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

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:

1. Count the total number of valence electrons (sum the group numbers; add electrons for anion charge, subtract for cation charge).
2. Draw a skeleton with the least electronegative central atom (H and F are never central; C is usually central in organic molecules).
3. Place a bonding pair between each pair of bonded atoms.
4. Distribute the remaining electrons as lone pairs to satisfy octets (start with the outer atoms).
5. 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:

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:

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:

1. The polar bonds are symmetrically arranged (linear, trigonal planar, tetrahedral with identical outer atoms, etc.).
2. The outer atoms are all the same.

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.

Examples in context

Example 1. Water treatment chlorination at Melbourne Water plants. Operators at the Winneke and Cardinia Creek treatment plants add chlorine gas $\text{Cl}_2$ to drinking water. The molecule is a homonuclear diatomic with a single covalent bond; electronegativity difference is zero, so the bond is non-polar and the molecule is non-polar. In water, $\text{Cl}_2$ disproportionates to $\text{HOCl}$ and $\text{HCl}$, both bent (V-shaped) molecules with two bonding pairs and lone pairs on oxygen or chlorine. $\text{HOCl}$ has a polar $\text{O-H}$ bond and a polar $\text{O-Cl}$ bond; the bent geometry means dipoles do not cancel, making $\text{HOCl}$ overall polar, which is why it dissolves readily in water and reaches pathogens efficiently.

Example 2. Greenhouse gases from Loy Yang power station. The brown-coal-fired Loy Yang generators release carbon dioxide and small amounts of methane. $\text{CO}_2$ is linear with two double bonds; each $\text{C=O}$ bond is polar (electronegativity difference about $1.0$), but the linear geometry means the two dipoles point in opposite directions and cancel, making the molecule non-polar overall. Methane $\text{CH}_4$ is tetrahedral with four $\text{C-H}$ bonds; the electronegativity difference is small ($0.4$) and the symmetric geometry cancels any small dipoles, so methane is essentially non-polar. Both gases absorb infrared radiation, however, because their bond vibrations create transient dipoles, which is why they act as greenhouse gases despite zero net dipole.

Try this

Q1. Predict the molecular shape and overall polarity of $\text{NH}_3$ and $\text{BF}_3$, and explain the difference. [3 marks]

• Cue. $\text{NH}_3$: trigonal pyramidal (3 bonding + 1 lone), polar overall. $\text{BF}_3$: trigonal planar (3 bonding, no lone), bond dipoles cancel, non-polar overall.

Q2. Consider $\text{SF}_4$ with four bonding pairs and one lone pair on sulfur. (a) Predict its geometry. (b) Calculate the approximate bond angles. (c) State whether the molecule is polar. [3 marks]

• Cue. (a) See-saw (distorted tetrahedron). (b) Axial $\text{F-S-F}$ near $173^{\circ}$, equatorial near $102^{\circ}$. (c) Polar; lone pair breaks symmetry.

Q3. Draw the Lewis structure of $\text{CO}_2$ and $\text{H}_2 \text{O}$. (a) State the geometry of each. (b) Compare their overall polarity. (c) Explain why $\text{CO}_2$ is a gas at $298 \, \text{K}$ while $\text{H}_2 \text{O}$ is a liquid. [2+2+2 marks]

• Cue. (a) $\text{CO}_2$ linear, $\text{H}_2 \text{O}$ bent. (b) $\text{CO}_2$ non-polar (cancellation); $\text{H}_2 \text{O}$ polar. (c) Polar water has hydrogen bonding; non-polar $\text{CO}_2$ only has weak dispersion forces.