Module 4 - 4.2: Drawing Lewis Structures for Molecular Compounds

.2 Drawing Lewis Structures for Molecular Compounds

The Lewis symbol for an atom is the atomic symbol surrounded by the correct number of valence

electrons. Lewis symbols for the second period are shown below. Lewis structures show how atoms

are bonded together.

There are standard conventions for drawing molecular compounds (i.e. compounds formed when two or

more atoms bond via a covalent bond).

• A bonding pair of electrons is represented by a line ( )

• A non-bonding pair of electrons (called a lone pair) is represented by a pair of dots (• • )

• Non-zero formal charges are always shown.

Examples:

Ammonia, NH3 has 3 bonding pairs of electrons and one lone pair of electrons. Carbon dioxide, CO2,

has 4 bonding pairs of electrons and four lone pairs of electrons.

Why is it important to draw proper Lewis structures?

Lewis structures show us how atoms bond to form molecules. We will soon see that we will be able to

correlate Lewis structures to VSEPR theory (Valence Shell Electron Pair Repulsion theory) to predict the

shapes of molecules, describe polarity and distribution of electron density. Molecular structures help us

understand and explain the physical properties and reactivity trends that chemists use for predicting how

molecules behave in chemical reactions and in different environments.

There are several different techniques that can be used to create Lewis Structures for molecular

compounds. We will briefly go over the traditional method for drawing Lewis structures and then

introduce you to a more expanded approach that helps you recognize structural and bonding patterns.

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# valence eCHEM 120/121 Module 4: Chemical Bonding and Structure page 4 of 22

Method 1: Traditional Approach for Drawing Lewis Structures

• Use group numbers to determine the # of valence e−’s

• Put a pair of e−’s between each pair of bonded atoms

• Distribute the remaining valence e−’s around atoms, start by filling terminal atom octets

first, but keep in mind the following:

(i) An H atom never has more than 2 valence e−’s

Therefore, the H atom never forms more than one bond and is always a “terminal” atom.

(ii) 2nd row atoms never have more than 8 valence e−’s

Also: C, N, O and F atoms almost never have less than 8 valence e−s around them.

(iii) Atoms from the 3rd period (and beyond) might have an “expanded” octet (more than 8

valence e−’s).

• Assign a formal charge to each of the atoms in your structure.

For the purposes of assigning formal charges, an atom “owns” all of its unshared electrons (i.e. lone

pairs) but only half of its bonding electrons. The formal charge concept is important in chemistry because

it helps us assess if a particular Lewis structure is important or reasonable.

Important notes:

• The assignment of formal charges is based on a set of rules for counting electrons. Since we cannot

create or destroy electrons, the sum of the formal charges must always equal the total charge

on the molecule or ion!!

• Place formal charges according to electronegativity, if possible (e.g. in the NCO− ion, it is better to

place the −1 formal charge on O rather than N because the O atom is more electronegative than

the N atom).

• Avoid placing like formal charges (e.g., +1 and +1 or −1 and −1) on adjacent atoms.

• All else being equal, the most important structure is (usually) the one with the smallest formal charges.

Almost always, formal charges do not exceed ±1. Structures having atoms with high formal charges (e.g.

+2, −2, +3, −3, etc.) are usually not important.

• In certain instances, it is sometimes possible to draw more than one acceptable Lewis structure for

a molecule. These structures are called resonance structures. We will learn more about

resonance structures in the next sections.

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