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B. Matter, Chemical Trends, and Chemical Bonding · Matter, Chemical Trends, and Chemical Bonding

Drawing Lewis Structures

B2.4 B3.4 A1.12About 75 minutes
Learning goals
  • I can count the valence electrons for a molecule or polyatomic ion.
  • I can draw Lewis structures for molecules with single, double and triple bonds.
  • I can draw Lewis structures for polyatomic ions and ionic compounds.
  • I can check a Lewis structure by counting electrons and octets.

Words to know

Lewis symbol
An element symbol with dots for its valence electrons.
Lewis structure
A diagram of a molecule or ion showing bonds as lines (shared pairs) and lone pairs as dots.
bonding pair
A pair of electrons shared between two atoms; one covalent bond.
lone pair
A pair of valence electrons on one atom that is not shared in a bond.
central atom
The atom in the middle of a molecule's skeleton, bonded to two or more other atoms.
polyatomic ion
A group of covalently bonded atoms that carries an overall charge, e.g. NH₄⁺ or OH⁻.

Drawing what you cannot see

In 1916 the American chemist Gilbert N. Lewis suggested a simple idea: a covalent bond is a pair of electrons shared between two atoms. To show it, he drew valence electrons as dots around element symbols. More than a hundred years later, chemists still sketch Lewis structures every day, because they show at a glance which atoms are bonded, how many bonds there are, and where the unshared electrons sit.

In Grade 9 you drew Lewis dot diagrams of atoms and of simple molecules. In this lesson you will learn a step-by-step method that works for molecules with double and triple bonds, for polyatomic ions, and for ionic compounds.

The pieces of a Lewis structure

  • Lewis symbol of an atom: the element symbol with one dot for each valence electron, placed on four sides, one at a time before pairing. Carbon has 4 dots, nitrogen 5, oxygen 6, chlorine 7.
  • Bonding pair: a shared pair of electrons, drawn as a line (–). A double bond (=) is 2 shared pairs and a triple bond (≡) is 3.
  • Lone pair (non-bonding pair): a pair of valence electrons that belongs to one atom and is not shared, drawn as two dots.

A structure with lines for bonds is often called a Lewis structure or structural formula with lone pairs. Every atom except hydrogen should end up with 8 electrons (count lone-pair electrons plus 2 for every bond line); hydrogen ends up with 2.

A method that always works (for this course)

  1. Count the total valence electrons of all the atoms. For a negative ion, add one electron per negative charge; for a positive ion, subtract one per positive charge.
  2. Draw the skeleton. The central atom is usually the one that forms the most bonds, usually the least electronegative atom (never hydrogen, which forms only one bond). Carbon is almost always central. Join every outer atom to the central atom with a single bond.
  3. Subtract 2 electrons for each bond you drew.
  4. Complete the octets of the outer atoms with lone pairs (hydrogen needs none).
  5. Put any electrons left over on the central atom as lone pairs.
  6. Check the central atom. If it has fewer than 8 electrons, move a lone pair from an outer atom into a shared position to make a double bond (or two moves for a triple bond). Repeat until all atoms have octets.
  7. Count again: the total number of electrons drawn must equal the number from step 1.
1. Count valence electrons2. Skeleton + single bonds3. Lone pairs to finish octetsHOH1 + 6 + 1 = 8 electrons(4 pairs)HOH2 bonds use 4 electrons4 left to placeHOHO: 2 bonds + 2 lone pairs = 8each H: 1 bond = 2
Building the Lewis structure of water. Oxygen ends up with two bonds and two lone pairs (8 electrons); each hydrogen has one bond (2 electrons).
Worked example

Example 1: Ammonia, NH₃

Draw the Lewis structure of NH₃.

  1. Step 1: N has 5, each H has 1: 5 + 3(1) = 8 electrons.
  2. Step 2: N is central (H can only form one bond). Draw three N–H single bonds.
  3. Step 3: 3 bonds use 6 electrons; 2 are left.
  4. Step 4: H atoms need no lone pairs.
  5. Step 5: put the last 2 electrons on N as one lone pair.
  6. Step 6: N has 3 bonds (6) + 1 lone pair (2) = 8. Each H has 2. Done.

H–N–H with a third H on N, and one lone pair on N. Ammonia has 3 bonding pairs and 1 lone pair.

Worked example

Example 2: Carbon dioxide, CO₂ (double bonds)

Draw the Lewis structure of CO₂.

  1. Step 1: C has 4, each O has 6: 4 + 2(6) = 16 electrons.
  2. Step 2: C is central (less electronegative than O): O–C–O. Two bonds use 4 electrons; 12 are left.
  3. Step 4: give each O three lone pairs (6 electrons each): 12 used, 0 left.
  4. Step 6: check C: only 2 bonds = 4 electrons. Not an octet.
  5. Move one lone pair from each O into a shared position with C. Now O=C=O.
  6. Check: C has 2 double bonds = 8 electrons. Each O has 1 double bond (4) + 2 lone pairs (4) = 8. Total drawn: 8 bonding + 8 lone-pair = 16 ✔.

O=C=O, with two lone pairs on each oxygen.

Worked example

Example 3: Hydrogen cyanide, HCN (triple bond)

Draw the Lewis structure of HCN.

  1. Step 1: 1 + 4 + 5 = 10 electrons.
  2. Step 2: C is central: H–C–N. 2 bonds use 4; 6 left.
  3. Step 4: complete N's octet with 3 lone pairs: 6 used, 0 left.
  4. Step 6: C has only 2 bonds (4 electrons). Move one lone pair from N: H–C=N. C now has 6. Move another: H–C≡N.
  5. Check: C has 1 single + 1 triple bond = 8. N has a triple bond (6) + 1 lone pair (2) = 8. H has 2.

H–C≡N with one lone pair on N.

Worked example

Example 4: The ammonium ion, NH₄⁺, and an ionic compound

Draw Lewis structures of (a) NH₄⁺ and (b) magnesium chloride.

  1. (a) Electrons: 5 + 4(1) − 1 (positive charge) = 8. N is central with four N–H bonds = 8 electrons. No lone pairs. Draw the structure in square brackets with the charge + outside.
  2. (b) Ionic compounds are not shared-pair structures. Mg (2 valence electrons) loses both: Mg²⁺ has no valence dots. Each Cl gains one: [Cl]⁻ with 8 dots in brackets.
  3. Write one ion for every ion in the formula: [Cl]⁻ [Mg]²⁺ [Cl]⁻ (each chloride with 8 dots).

Polyatomic ions and ions in ionic compounds are drawn in brackets with their charges. In ionic compounds, no bond lines are drawn between ions.

OCOCO₂ (16 e⁻): two double bondsNNN₂ (10 e⁻): triple bondHCNHCN (10 e⁻)CCHHHHC₂H₄ (12 e⁻): C=C double bondNHHHH+NH₄⁺ (8 e⁻)Na+Cl−NaCl: ionic, no shared pair
A gallery of Lewis structures: double bonds in CO₂ and C₂H₄, triple bonds in N₂ and HCN, a polyatomic ion (NH₄⁺) and an ionic compound (NaCl).

Quick checks

  • Carbon usually forms 4 bonds, nitrogen 3 (plus 1 lone pair), oxygen 2 (plus 2 lone pairs), hydrogen and the halogens 1 (halogens plus 3 lone pairs) in neutral molecules.
  • If your count of electrons drawn does not match step 1, you have an error.
  • A few molecules break the octet rule (for example BF₃, where boron has only 6 electrons, and SF₆, where sulfur has 12). You will study these exceptions in SCH4U; all the molecules in this course's practice follow the octet rule except where noted.
Just for fun

What did the oxygen say when it met magnesium?

"OMg!" (Fittingly, they form the ionic compound MgO.)

Hands-on activity

Lewis structures with kitchen counters

You need: a sheet of paper, a pencil, and about 30 small identical items to be electrons (dried beans, lentils or cereal pieces).

  1. Write the symbols C, O and O far apart on the paper. Count out the total valence electrons for CO₂ (16 beans).
  2. Place two beans between C and each O for single bonds. Place the remaining beans as lone pairs, outer atoms first.
  3. Count the beans around each atom (beans between two atoms count for both). Fix any atom with fewer than 8 by moving a lone pair into a bond position.
  4. Repeat for N₂ (10 beans), C₂H₄ (12 beans; the two carbons are joined to each other and each holds two H) and H₂O (8 beans).
  5. Draw each final structure on paper with lines for bonds.

What to look for: CO₂ ends with 4 beans between C and each O (two double bonds); N₂ ends with 6 beans between the atoms and 2 on each end; C₂H₄ has 4 beans between the carbons. In every case, the number of beans never changes, only their positions.

Your turn

1

How many valence electrons are in the Lewis structure of CH₄? Describe the structure.

2

Draw (describe) the Lewis structure of HCl.

3

How many valence electrons does the hydroxide ion, OH⁻, have? Describe its Lewis structure.

4

Draw (describe) the Lewis structure of N₂ and state the number of lone pairs.

5

Describe the Lewis structure of formaldehyde, CH₂O (C is central).

6

Describe the Lewis structures of the ions in calcium oxide.

More practice: warm-up

1

How many dots go around the Lewis symbol for sulfur?

2

What does a single line represent in a Lewis structure?

3

What is a lone pair?

4

Which atom is never central?

5

How many electrons does hydrogen have around it in a molecule?

More practice: on your own

1

Draw (describe) the Lewis structure of PCl₃.

2

Draw (describe) the Lewis structure of O₂.

3

Draw (describe) the Lewis structure of C₂H₂ (ethyne, H–C–C–H skeleton).

4

Find the total valence electrons in the carbonate ion, CO₃²⁻.

5

A student draws CO₂ as O–C–O with three lone pairs on each O and one lone pair on C. Find the error.

Challenge

1

Draw (describe) the Lewis structure of the cyanide ion, CN⁻.

2

Draw (describe) a Lewis structure for SO₂ that obeys the octet rule, and explain why chemists say a second, equally good structure exists.

3

Boron trifluoride, BF₃, has 24 valence electrons. Show that drawing three B–F single bonds and three lone pairs on each F uses all the electrons, and state what is unusual.

Check your understanding

Quiz

1.

How many valence electrons are shown in the Lewis structure of NH₃?

2.

Which molecule contains a triple bond?

3.

In the Lewis structure of water, oxygen has

4.

How many electrons are in the Lewis structure of the hydroxide ion, OH⁻?

5.

Which is the correct Lewis structure description for CO₂?

6.

In the Lewis structure of an ionic compound such as KCl,

For parents and teachers

Big idea: Expectation B2.4 asks students to draw Lewis structures to represent the bonds in ionic and molecular compounds. The step method in this lesson (count, skeleton, octets, fix with multiple bonds, recount) is the core skill.

How to help: Ask your teen to "count the beans" out loud. The total number of electrons never changes; only their positions do. Checking that the electrons drawn equal the electrons counted catches most mistakes.

Common difficulties:

  • Forgetting to add electrons for negative ions or subtract for positive ions.
  • Putting hydrogen in the middle.
  • Counting shared electrons for only one atom when checking octets.
  • Drawing bond lines between ions in ionic compounds. Ionic compounds are shown as separate bracketed ions.