U3Chem-3.5

Overview of Lewis Structures

• Lewis structures are representations of molecules that show how atoms are bonded and the distribution of electrons.

Key Concepts

Importance of Lewis Structures

• They help visualize the arrangement of electrons in a molecule and predict the shape, bond angles, and overall reactivity.

• Useful for understanding resonance, formal charges, and the stability of molecules.

Components of Lewis Structures

• Valence Electrons: The electrons in the outermost shell that participate in chemical bonding.

• Bonding Pairs: Pairs of electrons that are shared between atoms in covalent bonds. Represented as lines between two atoms in the structure.

• Lone Pairs: Non-bonded valence electrons that are not shared with other atoms but affect molecular shape and polarity.

Steps to Drawing Lewis Structures

1. Count the Valence Electrons: Add together all valence electrons from all atoms in the molecule.

   • Adjust for any charges: add electrons for negative charges and subtract for positive charges.

2. Determine the Central Atom: Usually, the least electronegative atom, which is usually not hydrogen.

3. Place Electrons: Draw bonds between the central atom and surrounding atoms, initially using single bonds.

4. Distribute Remaining Electrons: Add lone pairs to satisfy the octet rule (or duet rule for hydrogen).

5. Check for Octet Rule Satisfaction: Ensure all atoms (except for those that are satisfied with fewer than 8 electrons) have 8 electrons around them.

6. Adjust Structure if Necessary: If certain atoms do not have an octet, consider forming double or triple bonds by sharing lone pairs.

Common Examples

Water (H2O)

• Central atom: Oxygen

• Structure shows two single bonds between O and H, with two lone pairs on O.

Carbon Dioxide (CO2)

• Linear structure with double bonds between C and each O.

• No lone pairs on C, and each O has 4 electrons (2 lone pairs)

Ammonia (NH3)

• Central atom: Nitrogen

• Structure shows three single bonds with three H atoms and one lone pair on N.

Resonance Structures

• Some molecules cannot be represented by a single Lewis structure. Instead, multiple structures can represent the same molecule (e.g., ozone, benzene).

• Actual representation is a resonance hybrid that shows the delocalization of electrons.

Common Mistakes

• Not accounting for total valence electrons correctly.

• Forgetting the octet rule for certain elements (e.g., elements like phosphorus can exceed the octet).

• Underestimating the significance of lone pairs in determining molecular shape.

Additional Topics to Review

• Molecular Geometry: Understanding the shapes of molecules based on VSEPR theory.

• Polarity: Determining if molecules have polar or non-polar characters based on their Lewis structures.

• Formal Charges: Calculation to ensure the most stable structure with the lowest overall formal charge.

• Hybridization: Concept explaining the mixing of atomic orbitals to form new hybrid orbitals in molecules.