Chemistry · Ch 4 — Chemical Bonding and Molecular Structure

Lewis Representation of Simple Molecules (the Lewis Structures)

4.1.3

Lewis Representation of Simple Molecules (the Lewis Structures)

The Lewis dot structure is a symbolic picture that shows how atoms in a molecule share their valence electrons. It is built around two core ideas: that bonding involves the sharing of electron pairs, and that atoms tend to achieve a stable noble-gas electron configuration — the octet rule (or duplet rule for hydrogen). While this representation does not capture every subtlety of molecular behaviour, it gives a remarkably useful first picture of how a molecule forms and why it has certain properties.

Writing a Lewis structure is a systematic process. The textbook lays out a clear sequence of steps, and we will follow them exactly.

Step 1: Count the total number of valence electrons

The first thing you need is the total number of electrons available for bonding. This is simply the sum of the valence electrons of all the atoms in the molecule or ion.

For a neutral molecule like methane ( $\mathrm{CH_4}$ ):

Carbon has 4 valence electrons ( $2s^2 2p^2$ ).
Each hydrogen has 1 valence electron ( $1s^1$ ), and there are four hydrogens.
Total = $4 + (4 \times 1) = 8$ valence electrons.

For an ion, you must adjust for the charge:

For an anion (negative charge), add one electron for each negative charge.
For a cation (positive charge), subtract one electron for each positive charge.

Consider the carbonate ion, $\mathrm{CO_3^{2-}}$ :

Carbon: 4 valence electrons.
Each oxygen: 6 valence electrons, and there are three oxygens.
The $2-$ charge means two extra electrons.
Total = $4 + (3 \times 6) + 2 = 24$ valence electrons.

Consider the ammonium ion, $\mathrm{NH_4^+}$ :

Nitrogen: 5 valence electrons.
Each hydrogen: 1 valence electron, and there are four hydrogens.
The $1+$ charge means one electron is lost.
Total = $5 + (4 \times 1) - 1 = 8$ valence electrons.

Watch out

A common mistake is to forget the charge adjustment. Always check: a negative charge means more electrons; a positive charge means fewer.

Step 2: Write the skeletal structure

You need to know (or guess intelligently) which atoms are bonded to which. The chemical symbols are arranged in a rough framework. For simple molecules, this is often obvious from the formula.

Step 3: Place the least electronegative atom at the centre

In general, the atom with the lowest electronegativity occupies the central position. The more electronegative atoms (like fluorine, oxygen, chlorine) are placed at the terminal (outer) positions.

For example:

In $\mathrm{NF_3}$ , nitrogen is less electronegative than fluorine, so nitrogen is the central atom.
In $\mathrm{CO_3^{2-}}$ , carbon is less electronegative than oxygen, so carbon is central.

Tip

Hydrogen is never the central atom (except in $\mathrm{H_2}$ itself). It can only form one bond, so it always goes on the outside.

Step 4: Distribute electrons as bonding pairs

First, connect the central atom to each terminal atom with a single bond (one shared pair of electrons). Count how many electrons you have used for these single bonds. Each single bond uses 2 electrons.

Step 5: Complete octets on terminal atoms (except hydrogen)

Place the remaining electrons around the terminal atoms (starting with the most electronegative) as lone pairs, so that each terminal atom (except hydrogen) gets an octet. Hydrogen gets only a duplet (2 electrons).

Step 6: Place any leftover electrons on the central atom

If there are electrons left after completing the octets of the terminal atoms, place them as lone pairs on the central atom.

Step 7: If the central atom does not have an octet, form multiple bonds

This is the critical step. If the central atom still has fewer than 8 electrons after step 6, you must convert one or more lone pairs from a terminal atom into bonding pairs — that is, form a double or triple bond between the central atom and that terminal atom. The goal is to give every bonded atom (except hydrogen) a complete octet.

Problems 4.1 (the Lewis dot structure of CO) and 4.2 (the Lewis structure of the nitrite ion, $\mathrm{NO_2^-}$ ) in this section's Problems group below walk these steps through in full.

Table 4.1: Lewis Representations of Some Molecules

Table 4.1The Lewis Representation of Some Molecules

Molecule/Ion	Lewis (bond-line) representation
$H_2$	$H-H$ (each H attains the He duplet)
$O_2$	$\ddot{O}=\ddot{O}$ (a double bond + two lone pairs on each O)
$O_3$	$\ddot{O}-\overset{+}{O}=\ddot{O}$ (bent; one single + one double bond)
$NF_3$	N single-bonded to three F, with one lone pair on N and three on each F

The textbook provides a table of Lewis structures for common molecules. Here are the key examples, with the central atom identified and the electron count verified.

Molecule/Ion	Total Valence Electrons	Central Atom	Lewis Structure (Description)
$\mathrm{CH_4}$	$4 + (4 \times 1) = 8$	Carbon	Carbon forms four single bonds to four hydrogens. No lone pairs. Each H has a duplet.
$\mathrm{NH_3}$	$5 + (3 \times 1) = 8$	Nitrogen	Nitrogen forms three single bonds to three hydrogens and has one lone pair.

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