2.5 Lewis Diagrams

Lewis Diagram AP Chemistry Summary

Lewis diagrams show how valence electrons are arranged in a molecule using dots for lone pairs and lines for bonds. You draw them by counting total valence electrons, connecting atoms, completing octets, and adjusting with multiple bonds when needed. Learning this skill is essential because it sets up molecular shape, polarity, and bonding predictions later in AP Chemistry.

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Why This Matters for the AP Chemistry Exam

On the AP Chemistry exam, you are expected to construct Lewis diagrams and then use them to make claims about a molecule's structure and behavior. A correct diagram is the starting point for predicting shape, bond angles, polarity, and reactivity, so errors here ripple into everything that follows in Unit 2.

This skill shows up in both multiple-choice questions and free-response problems. You may be asked to draw a structure, identify lone pairs and bonding pairs, count valence electrons, or explain why one arrangement is more reasonable than another. Getting the electron count and octets right is what makes the rest of your reasoning hold up.

Key Takeaways

• Count total valence electrons first, including extra electrons for negative ions and fewer for positive ions.
• The least electronegative atom usually goes in the center, and hydrogen always stays terminal.
• Lone pairs sit on a single atom; bonding pairs are shared and shown as lines.
• If you have too many electrons after completing octets, form double or triple bonds.
• Know the octet exceptions: hydrogen (2 electrons), small atoms like Be and B, odd-electron species, and central atoms that can hold more than 8.
• For ions, draw brackets and place the overall charge outside.

What Lewis Diagrams Show

Lewis diagrams, also called Lewis structures, are a way to draw molecules so you can see the valence electrons and the bonds between atoms. Being able to draw any molecule opens the door to predicting its shape and how it reacts, which is why this is one of the most useful skills in the course.

Most Lewis diagrams follow the octet rule, the idea that atoms tend to bond so each one ends up with eight valence electrons (a full outer shell). Hydrogen is the main exception because it is full with just two electrons.

In a Lewis structure, valence electrons fall into two groups:

1. Lone pairs: pairs of electrons that stay on a single atom and are not shared.
2. Bonding pairs: pairs of electrons shared between atoms, usually drawn as a line (dash).

Drawing Ionic Lewis Structures

Ionic structures are simpler because ionic bonds form when electrons transfer from a metal to a nonmetal rather than being shared.

1. Write the empirical formula so you know which elements are present and how many of each.
2. Identify the valence electrons on the metal and nonmetal.
3. Transfer valence electrons from the metal to the nonmetal until both reach a full octet.
4. Draw brackets around each ion and write its charge outside the brackets.

Example: NaBr (Sodium Bromide)

• The formula NaBr tells you there is one atom of each element.
• Na has 1 valence electron and Br has 7, which you can check by their groups on the periodic table.
• One electron transfers from Na to Br, giving both a full octet. This creates a sodium cation and a bromide anion, so each charge goes outside its bracket.

Drawing Covalent Lewis Structures

Covalent structures take more steps because electrons are shared instead of transferred.

1. Count the total number of valence electrons from the formula.
2. Choose the central atom. It is usually the least electronegative atom, and hydrogen is never central.
3. Place the outer atoms around the center and connect them with single bonds.
4. Complete octets on the outer atoms (and the center where possible).
5. Count the electrons you have used. If you have too many, remove lone pairs from adjacent atoms and form double or triple bonds. If you have too few, see the exceptions below.

Exceptions to the Octet Rule

Not every atom follows the octet rule, so watch for these cases:

• Some atoms hold fewer than eight electrons. Hydrogen maxes out at 2, beryllium at 4, and boron at 6.
• A central atom can hold more than eight electrons only if its atomic number is 14 or greater. Add extra electrons there when your count requires it.
• Odd-electron species (like NO) cannot pair every electron into neat octets, so one electron is left unpaired.
• When electrons run short, atoms often form multiple bonds. For example, CO2 has carbon double-bonded to each oxygen.

Worked Examples

1) Lewis structure for O2

• Oxygen is in group 16 with 6 valence electrons, so two oxygens give 12 total.
• Place the two oxygen atoms side by side (no real central atom here) and connect them with a single bond.
• Add 3 lone pairs to each oxygen to complete octets. This uses 14 electrons.
• You have 2 too many, so convert the single bond into a double bond by removing one lone pair from each atom.
• Recount: 12 electrons total with a full octet on each oxygen. That is the correct structure.

2) Lewis diagram for CS2

• Carbon has 4 valence electrons and each sulfur has 6, giving 16 total.
• Carbon is the central atom.
• Connect each sulfur to carbon with a single bond, then complete octets on all three atoms. This overshoots to 20 electrons.
• You need 16, so form two double bonds (one to each sulfur) to reach the correct count.

3) Lewis diagram for XeF2

• Xe has 8 valence electrons and each F has 7, giving 22 total.
• Xe is the central atom.
• Connect each fluorine with a single bond and complete octets on the fluorines and xenon. That uses 20 electrons.
• You still need 2 more. Since xenon's atomic number is 54 (well above 14), it can hold extra electrons, so place the remaining pair on xenon.

When a central atom carries three lone pairs and two bonds, keep the lone pairs in pairs of two. You cannot split them into groups of three electrons on each side. That arrangement would be incorrect.

Answers to Practice Structures

Magnesium Chloride

Write the empirical formula first and make sure the charges cancel to zero. Mg is +2 and Cl is -1, so the formula is MgCl2 with two chlorine atoms. You can also draw one Cl with a 2 written outside the bracket to show two chloride ions. Do not confuse a coefficient with a subscript or superscript.

N2

After the steps, you find 14 electrons, but you only need 10 (each nitrogen has 5). A double bond still leaves you short, so nitrogen forms a triple bond to reach the correct count and full octets.

NH4+

A positive charge means one electron is missing, while a negative charge means electrons are added. NH4+ would look like it has 9 valence electrons, but the +1 charge drops it to 8 total. Draw the structure with brackets and the charge outside.

How to Use This on the AP Chemistry Exam

Free Response

When a question asks you to draw a Lewis structure, count valence electrons carefully and adjust for the charge on any ion. Show lone pairs and bonds clearly, since graders look for the correct electron count and proper octets. After drawing, you may need to use the structure to justify a claim about shape or polarity, so make the diagram accurate before you reason from it.

Problem Solving

Use a consistent process: total the valence electrons, place the central atom, connect with single bonds, complete octets, then form multiple bonds if you have leftover or missing electrons. This routine keeps you from skipping the electron count, which is the most common source of errors.

Common Trap

Confusing molecular geometry with the Lewis diagram itself trips up many students. The Lewis structure shows electron arrangement on paper; it does not directly show the 3D shape. You will use it to predict shape later, but the two are not the same representation.

Common Misconceptions

• A Lewis diagram is not a picture of the molecule's real 3D shape. It shows electron placement, and you use it to predict geometry separately.
• The octet rule is not absolute. Hydrogen, small atoms like Be and B, odd-electron species, and larger central atoms can break it.
• Forgetting to adjust for ionic charge is a frequent mistake. Add electrons for negative ions and remove them for positive ions before drawing.
• A single line (dash) represents a shared pair of two electrons, not one electron. Miscounting bonds as single electrons throws off the whole structure.
• The central atom is usually the least electronegative atom, not the first one written in the formula, and hydrogen is never central.

Related AP Chemistry Guides

• 2.1 Types of Chemical Bonds
• Unit 2 Overview: Compound Structure and Properties
• 2.2 Intramolecular Force and Potential Energy
• 2.3 Structure of Ionic Solids
• 2.4 Structure of Metals and Alloys
• 2.7 VSEPR and Bond Hybridization