16 - Lewis Theory

Understanding Lewis Theory

  • Key to understanding bonding in atoms.

  • Requires knowledge of electron configuration and valence electrons.

  • Having a periodic table handy is useful.

Valence Electrons

  • Valence electrons are the outer shell electrons, important for bonding.

  • Example:

    • Oxygen: Electron configuration is 1s² 2s² 2p⁴ → 6 valence electrons

    • Sulfur: Electron configuration is 1s² 2s² 2p⁶ 3s² 3p⁴ → 6 valence electrons

    • Selenium: Electron configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁴ → 6 valence electrons

  • Valence electrons can be quickly identified using the periodic table (groups).

Octet Rule

  • Atoms strive for 8 valence electrons to achieve stability (like Noble gases).

  • Example: Nitrogen, with 5 valence electrons, needs 3 more to complete its octet.

Electronegativity

  • Electronegativity: the ability of an atom to attract bonding electrons.

  • Trends in the periodic table:

    • Fluorine is most electronegative.

    • Electronegativity tends to increase from left to right and bottom to top of the periodic table.

  • Affects atom size; more electronegative elements tend to be smaller.

Lewis Structures

  • Steps to create Lewis structures:

    1. Count total valence electrons.

    2. Draw the skeleton structure (least electronegative in the center).

    3. Connect atoms with single bonds (each bond = 2 electrons).

    4. Distribute remaining valence electrons to fulfill octet rule, starting from outer atoms.

    5. If central atom does not fulfill octet, create double or triple bonds as necessary.

Examples of Lewis Structures

  • PCl₃:

    • Valence electrons: P (5) + 3Cl (7 each) = 26.

  • CH₂O:

    • Valence electrons: C (4) + H (1 each) + O (6) = 12.

    • Final structure may require double bond between C and O to fulfill octet.

Formal Charge

  • Formal charge is calculated by:

    • Formal charge = valence electrons - (non-bonded electrons + 1/2 bonded electrons).

  • Examples:

    • NH₄⁺: N has +1 charge, all H's have a charge of 0.

    • For CN⁻: C gets -1 charge with resonance structures achievable for stability.

Resonance

  • Situations where multiple valid Lewis structures exist for a molecule.

  • Example: CN⁻ can resonate between multiple forms.

  • Stability is assessed based on formal charges: generally, the structure with formal charges closest to zero is more stable.

Exceptions to Octet Rule

  • Some elements can exist with less than an octet (e.g., BF₃).

  • Some halogens cannot expand their valence shell (e.g., no double bonds with Fluorine).

Conclusion

  • Familiarity with these concepts involves significant practice.

  • Reinforce your understanding with multiple examples and exercises.