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Lewis Dot Structures
Electron Dots, Octet Rule, Single/Double/Triple Bonds
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Lewis dot structures are simple diagrams that show how valence electrons are arranged around atoms in molecules and ions. They help students predict bonding patterns, molecular shapes, and chemical reactivity. By focusing on outer electrons, these structures connect atomic structure to the way substances actually form and behave. They are a core tool in chemistry because they make invisible electron arrangements easier to visualize.
To draw a Lewis structure, you count total valence electrons, choose a central atom, connect atoms with single bonds, and then distribute the remaining electrons to satisfy octets when possible. If needed, you form double or triple bonds to give atoms the correct number of electrons. Lewis structures also help identify lone pairs, bonding pairs, and formal charges. A molecule like CO2 shows this clearly, with carbon in the center and double bonds to each oxygen.
Key Facts
- Valence electrons are the outermost electrons and are the ones shown in Lewis dot structures.
- A single bond represents 2 shared electrons, a double bond represents 4, and a triple bond represents 6.
- Total valence electrons in a molecule = sum of the valence electrons from all atoms, adjusted for charge.
- Formal charge = valence electrons - nonbonding electrons - 1/2(bonding electrons).
- Most main group atoms follow the octet rule and tend to have 8 electrons around them in a stable structure.
- For CO2, total valence electrons = 4 + 2(6) = 16, and the best Lewis structure is O=C=O.
Vocabulary
- Valence electron
- An electron in the outer energy level of an atom that can participate in bonding.
- Lewis dot structure
- A diagram that uses dots and lines to show valence electrons and bonds in a molecule or ion.
- Lone pair
- A pair of valence electrons on an atom that is not shared in a bond.
- Octet rule
- The guideline that many atoms are most stable when they have 8 electrons in their valence shell.
- Formal charge
- The charge assigned to an atom in a Lewis structure based on how electrons are distributed.
Common Mistakes to Avoid
- Forgetting to count the total valence electrons first, which leads to structures with too many or too few electrons. Always add electrons from every atom before drawing bonds.
- Choosing hydrogen as the central atom, which is wrong because hydrogen forms only one bond and cannot hold more than 2 electrons. The central atom is usually the least electronegative atom that can form multiple bonds.
- Leaving second period atoms like carbon, nitrogen, and oxygen with incomplete octets when enough electrons are available. These atoms usually need 8 electrons in correct Lewis structures.
- Ignoring the charge on an ion, which changes the total electron count. Add one electron for each negative charge and subtract one for each positive charge.
Practice Questions
- 1 Draw the Lewis dot structure for H2O. How many valence electrons are there in total, how many bonds are present, and how many lone pairs are on oxygen?
- 2 Draw the Lewis dot structure for NH3. Count the total valence electrons and determine the number of bonding pairs and lone pairs around nitrogen.
- 3 CO2 is drawn as O=C=O instead of O-C-O with only single bonds. Explain why the double bond structure is preferred in terms of octets and formal charges.