Embedded Files
Introduction to Lewis Structures
Lewis structures are simplified visualizations of chemical bonding between atoms. They help chemists predict the molecular geometry, reactivity, and polarity of molecules.
Concepts to Know:
Covalent Bonding: Formed when valence electrons are shared between two atoms.
Each bond consists of two electrons, one contributed by each atom.
There are single bonds, double bonds, and triple bonds.
Duet Rule: Applies to hydrogen (H) and helium (He), where a full valence shell consists of two electrons from their only electron orbital.
Octet Rule: Applies to most elements with an outermost p-subshell, where a stable configuration has eight valence electrons.
Incomplete Octet: Some elements are stable with less than 8 electrons. For example, boron (B) only needs 6 to achieve stability, while beryllium (Be) only needs 4.
Radicals: Neutral molecules that contain unpaired electrons, making them highly reactive. Radicals often have common ion counterparts that are more stable.
To identify radicals while drawing Lewis structures, check to see if the calculated valence electron count is an odd number. The unpaired lone electron is typically located on the most electronegative atom in the molecule.
Hypervalency: Atoms capable of hypervalency (elements in period 3 and below) don't follow the octet rule when covalently bonding. This is because they have access to d-orbitals for bonding. These atoms can form bonds equal to their valence electron count, allowing them to exceed the octet.
Valence electrons for main group elements are based on the periodic table:
Group 1: 1 electron.
Group 2: 2 electrons.
Group 13 (3A): 3 electrons.
Group 14 (4A): 4 electrons.
Group 15 (5A): 5 electrons.
Group 16 (6A): 6 electrons.
Group 17 (7A): 7 electrons.
Group 18 (8A): 8 electrons.
Excluding Helium (He).
To view the tutorial for VSEPR Theory and Molecular Shapes, click HERE.
How to Draw Lewis Structures (With Example Pictures)
Step 1: Count the Valence (Outermost) Electrons of Atoms
Count valence electrons of all atoms in the molecule/ion.
For polyatomic ions, add or subtract up the extra electrons based on their charges.
Add electrons for negative charges.
Subtract electrons for positive charges.
Step 2: Determine the Central Atom
The atom that could form the most amount of bonds is usually the central atom of a molecule.
Typically the least electronegative element with p-subshells and/or has a valency closest to 4.
Atoms capable of hypervalency can also serve as central atoms.
Hydrogen is always terminal because it could only form one bond (duet rule).
Atoms with a valency closest to 7 (excluding noble gases) usually end up as terminal atoms, as they typically form less bonds than the central atom.
Step 3: Connect the Atoms via Bonds and Complete
First, connect all atoms together with single bonds without exceeding their valence limit. Each bond consists of two electrons, with one electron originating from each atom.
Then, make up for atoms with incomplete octets by adding more bonds.
The maximum bond count between atoms is 3 (triple bond).
If an atom cannot make more bonds, add lone pairs (two electrons each, so draw two dots next to each other).
If multiple valid structures are possible, resonance structures can be drawn to represent the molecule's true bonding.
Double check the electron count. Again, every bond and lone pair consists of two electrons. The valence electrons should match what was calculated in step 1.
Examples of Lewis Structures
Resonance Structures
Some molecules or ions can be represented by more than one valid Lewis structure. These groups of structures, known as resonance structures, show the variations in how electrons could be distributed within the molecule. The actual structure of the molecule is not any one of the individual resonance structures, but rather a combination of all possible resonance structures, called the resonance hybrid.
Criteria for molecules to have resonance:
Must have the presence of different bond counts and different lone pair counts, or there would be no variation otherwise.
This difference can be moved across the molecule.
The molecule must have electrons that can be delocalized across the molecule.
Often caused by π-bonds.
All resonance structures must be equal in stability.
Characteristics of molecules with resonance:
Due to the delocalization of electrons, all bonds between atoms in a resonance molecule appear to be equal in length, as the electron densities are averaged over the different resonance forms and evenly distributed across the molecule.
Resonance Structures
Some molecules or ions can be represented by more than one valid Lewis structure. These groups of structures, known as resonance structures, altogether create a resonance hybrid, which is a combination of all possible resonance structures.
Criteria for molecules to have resonance:
Must have the same arrangement of atoms.
Must have multiple ways to draw Lewis structure bonds in the molecule.
Must have pi-bonds that delocalize.
Characteristics of molecules with resonance:
Due to the delocalization of electrons, all bonds between atoms in a resonance molecule appear to be equal in length, as the electron densities are averaged over the different resonance forms and evenly distributed across the molecule.
Molecules with resonance are more stable due to the distribution of charges.
Example - Resonance Structures of NO₃⁻
Page updated
Google Sites
Report abuse