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B2.4 · Draw Lewis structures for ionic and molecular compounds

Learn to draw lewis structures for ionic and molecular compounds through clear examples and targeted practice.

Ontario Grade 11 Chemistry

Matter, Chemical Trends, and Chemical Bonding

Ontario Grade 11 Chemistry · B2.4

Some substances conduct electricity when dissolved in water, while others do not. This observation can prompt questions about how their particles differ, but it does not show the particles directly. A Lewis structure is a model that shows selected outer electrons. In an ionic compound, it shows ions formed when electrons transfer from one atom to another. In a molecular compound, it shows atoms sharing pairs of electrons. You will learn to draw and check both kinds of structures.

What you will learn

Explain how Lewis structures represent valence electrons.
Draw ionic Lewis structures with the correct ions, brackets, charges, and electron dots.
Draw molecular Lewis structures with shared pairs and lone pairs.
Check a structure against its formula, electron count, and charge.

1. Review: atoms and valence electrons

An atom has a small, positively charged nucleus and electrons around it. Valence electrons are the electrons in an atom’s outer occupied shell. Lewis structures show these electrons as dots or as lines between bonded atoms.

For the main-group elements used here, the periodic table helps you find the number of valence electrons. Magnesium is in Group 2 and has two. Chlorine is in Group 17 and has seven. Hydrogen has one, and oxygen has six.

A Lewis symbol is an element symbol with dots around it to show its valence electrons. Begin by placing one dot on each side before pairing dots on a side. The starting side does not matter. The number of dots does.

These dots are a model for keeping track of selected electrons. They are not a picture of an atom’s size or the paths followed by its electrons.

Lewis structures represent valence electrons.
Use the periodic table to find valence-electron counts for these main-group elements.
A Lewis structure is a model, not a picture of an atom’s size or electron paths.

2. Ionic compounds: show ions and transferred electrons

An ion is an atom or group of atoms with a net charge because electrons have been lost or gained. An atom that loses electrons becomes a positive ion. An atom that gains electrons becomes a negative ion.

For the common main-group ions in these examples, metals lose valence electrons and non-metals gain electrons. This gives the ions a full outer shell, usually with eight electrons. A metal ion that has lost its valence electrons is often shown without dots. Show eight dots around a non-metal ion that has gained electrons.

In an ionic Lewis structure, put each ion in its own brackets and write its charge outside the brackets. Brackets and charges show that the drawing represents ions rather than neutral atoms.

Use the compound’s formula to decide how many ions to draw. For example, magnesium chloride contains one magnesium ion for every two chloride ions. Magnesium loses two electrons; each chlorine gains one. The electron transfer gives the correct ion charges, and the charges must add to zero overall. Do not change the formula to make the drawing look balanced.

Draw each ion in brackets and put its charge outside.
Show eight valence-electron dots around each non-metal anion in these examples.
Match the ion counts to the formula and check that the total charge is zero.

3. Molecular compounds: show shared and unshared pairs

A molecular compound contains atoms joined by covalent bonds. A covalent bond is a pair of electrons shared between atoms. Draw a shared pair as two dots between atoms or as one line. One line represents one shared pair.

A lone pair is a pair of valence electrons that is not shared in a bond. Draw lone pairs as pairs of dots beside an atom. In the simple structures in this lesson, hydrogen has two electrons around it, while the other common atoms shown have eight.

To draw a molecular structure, first count the valence electrons from all atoms in the formula. Connect the atoms with shared pairs. Hydrogen can make only one bond, so it goes at an outside position. Add lone pairs to complete the usual outer-shell pattern, then count the electrons shown.

Do not use brackets and ion charges for each atom in a molecular structure. The shared pairs show how the atoms are connected. When checking the electron total, count a bond line as one pair, not as extra electrons in addition to that pair.

A single bond line represents one shared pair of electrons.
Lone pairs are not shared in bonds.
Check the total electron count and the outer-shell pattern.

4. Choose a drawing style and check it

Start with the formula. For the compounds in this lesson, a metal combined with a non-metal is represented as an ionic compound. A compound made of non-metals is represented as a molecular compound.

For an ionic structure, draw the number of bracketed ions shown by the formula. Include their charges and the anion electron dots. Confirm that the charges add to zero. For a molecular structure, count all valence electrons, draw shared pairs and lone pairs, and confirm the total and outer-shell pattern.

The two drawing styles communicate different particle models. Ionic structures show ions and their charges. Molecular structures show atoms connected by shared pairs. In both cases, the formula and electron count provide useful checks.

Decide whether the compound is ionic or molecular before drawing.
For an ionic compound, check ion count, dots, brackets, and total charge.
For a molecular compound, check bonds, lone pairs, electron total, and outer shells.

Worked example

Draw magnesium chloride and water

Draw Lewis structures for magnesium chloride and water. Use brackets for ions and lines for shared pairs where appropriate.

Classify the compounds
Magnesium is a metal and chlorine is a non-metal, so represent magnesium chloride as ions. Hydrogen and oxygen are non-metals, so represent water as a molecule with shared electron pairs.
Draw magnesium chloride
Magnesium loses two valence electrons and becomes a $2+$ ion. Each chlorine atom gains one electron and becomes a $1-$ ion. The formula requires one magnesium ion and two chloride ions. Show eight dots around each chloride ion, including the electron it gained, and no dots around the magnesium ion.
$[\mathrm{Mg}]^{2+}\quad[\mathrm{Cl}(\cdot\cdot)_4]^{-}\quad[\mathrm{Cl}(\cdot\cdot)_4]^{-}$
Draw water
Water has eight valence electrons in total: six from oxygen and one from each hydrogen. Oxygen is in the centre. It contributes one electron to each oxygen–hydrogen bond, forming two shared pairs. Oxygen retains four of its own electrons as two lone pairs. Count each shared pair once in the total: the bonds show four electrons, and the lone pairs show four more. Each hydrogen has two electrons around it, and oxygen has eight when its bonds and lone pairs are counted.
$\begin{array}{c}\cdot\cdot\\[-2pt]\mathrm{H}-\mathrm{O}-\mathrm{H}\\[-2pt]\cdot\cdot\end{array}$
Check the ions and electron count
The magnesium chloride ion counts and charges match its formula, and the total charge is zero. In water, two bonds show four electrons and two lone pairs show four more, for eight valence electrons altogether.
$1(+2)+2(-1)=0$

Answer: Magnesium chloride is drawn as one bracketed magnesium ion and two bracketed chloride ions, each with eight dots. Water is drawn with two single bonds to oxygen and two lone pairs on oxygen.

Check: Magnesium chloride has a neutral total charge. The water structure shows eight valence electrons, with two around each hydrogen and eight around oxygen.

Common mistakes and how to avoid them

Drawing an ionic compound with shared bonds between its ions.

Correction: Show separate bracketed ions with charges. Use shared pairs for molecular compounds.

Leaving the gained electrons off an anion.

Correction: Show eight dots around each non-metal anion in these examples, including the electrons gained.

Leaving charges off bracketed ions.

Correction: Write each ion’s charge outside its bracket and check that the compound’s total charge is zero.

Counting a bond line and then counting its electron pair again.

Correction: A line represents one shared pair. Count that pair once.

Giving hydrogen eight electrons in a molecular structure.

Correction: Hydrogen follows the two-electron pattern in these structures.

Lesson summary

Lewis structures show valence electrons as dots or bond lines.
Ionic structures use brackets and charges to show ions formed by electron transfer.
Molecular structures use shared pairs for bonds and dots for lone pairs.
Check the formula, ion charges, electron count, and outer-shell pattern.

Check your understanding

Question 1

Which description best fits an ionic Lewis structure?

Separate bracketed ions with charges
Atoms connected only by shared-pair lines
One bracket around the neutral formula with no charges
Element symbols with no dots or charges

Show answer and explanation

Separate bracketed ions with charges

An ionic Lewis structure shows the separate ions, their brackets, and their charges.

Question 2

How many lone pairs are shown on oxygen in a simple Lewis structure for water?

One
Two
Three
Four

Show answer and explanation

Two

Oxygen contributes one electron to each of its two bonds to hydrogen. Its other four valence electrons remain as two lone pairs.

Question 3

What does one line between two atoms represent in a molecular Lewis structure?

One shared pair of electrons
One positive charge
Two lone pairs
A pair of ions

Show answer and explanation

One shared pair of electrons

A single bond line represents one shared electron pair.

Key terms

Valence electrons: Electrons in an atom’s outer occupied shell, shown in a Lewis structure.
Lewis structure: A drawing that represents valence electrons, bonds, and sometimes ion charges.
Ion: An atom or group of atoms with a net charge because electrons have been lost or gained.
Covalent bond: A shared pair of electrons between atoms.
Lone pair: A pair of valence electrons that is not shared in a bond.

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About this lesson and its review

Published by DoAssignment. This reviewed lesson follows Ontario Grade 11 Chemistry (SCH3U), expectation B2.4. It is a study resource, not an official curriculum publication.

Before publication, content is checked for structure, mathematical or chemical notation, calculations, course boundaries, and readability. Errors can still occur, so corrections are welcomed.

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