General Chemistry: Lewis Structures, Resonance, and Formal Charge - 30

General Chemistry for Engineers - CHEM 1201 Lecture Notes

Introduction

  • Instructor: Carolyn Kohlmeier

  • Department: Chemical and Biological Engineering

  • Lecture Information: Week 11, Lecture 30

  • Covered Topics:

    • Last time's topics:

    1. Lewis structures

    2. Covalent bonding

    3. Electronegativity

    • Today's topics:

    1. Resonance

    2. Formal charge

Lewis Structures: Review

  1. Write the correct skeletal structure.

  2. Calculate the total number of electrons.

  3. Distribute the electrons to give octets or duets where possible.

  4. Use double/triple bonds to satisfy the octet rule.

    • Examples:

      • Carbon dioxide (CO₂)

      • Ammonia (NH₃)

  • Placement of atoms:

    • Place hydrogen (H) and high-electronegativity (EN) atoms on the outside.

  • Example Calculation for Electron Count:

    • For CO₂:

    • 4 (for C) + 2·6 (for O) = 16 electrons total

    • For NH₃:

    • 5 (for N) + 3·1 (for H) = 8 electrons total

Lewis Structures: Visual Representation

  • Example diagrams (not shown) include various electron dot diagrams

  • Several structures are judged based on their correctness.

Structuring Lewis Structures for CH₂Cl₂

  • Presented options A, B, C, D, which depict various arrangements.

    • Correct arrangement must adhere to electron count.

  • Calculation Example:

    • Total Valence Electrons for CH₂Cl₂:

    • Distribution must account for double bonds (if any).

Example Problem for Lewis Structure

  • Draw a Lewis structure for COCl₂.

  • Question: How many valence electrons are there?

    • Options:

      • (A) 20

      • (B) 22

      • (C) 24

      • (D) 26

      • (E) 28

Double Bonds in Lewis Structures

  • Draw a Lewis structure for COCl₂.

  • Question: How many double bonds are present?

    • Options:

      • (A) 0

      • (B) 1

      • (C) 2

      • (D) 3

Resonance Structures

  • Resonance structure:

    • A representation of a molecule that can be depicted in two or more ways while retaining the same skeletal formula but differing in electron arrangements.

  • Resonance hybrid:

    • The actual structure is an average of all possible resonance structures.

  • Formal charge (FC):

    • Represents the hypothetical charge of an atom if all electrons were shared equally.

    • Calculated using:

    • FC=(valenceextelectrons)−(non−bondingextelectrons)−0.5imes(bondedextelectrons)FC = (valence ext{ electrons}) - (non-bonding ext{ electrons}) - 0.5 imes (bonded ext{ electrons})

  • Guidelines for evaluating formal charge:

    1. The sum of formal charges must equal the charge of the molecule.

    2. Smaller formal charges are preferred over larger formal charges.

    3. Negative formal charges should reside on more electronegative atoms when possible.

  • Example: O₃ (Ozone)

    • Electron calculation:

      • 3·6 = 18 electrons total

    • Comparison of structures based on formal charge shows Structure A as preferable due to lower charges.

Assignment of Formal Charge

  • Helpful in determining the most appropriate Lewis structure under conditions of multiple possibilities.

  • Reiterate how to calculate formal charge for N₂O:

    • Using the formula:

    • FC=(valenceexte−)−(non−bondingexte−)−rac(bondedexte−)2FC = (valence ext{ e}⁻) - (non-bonding ext{ e}⁻) - rac{(bonded ext{ e}⁻)}{2}

Example Structure for NO₃

  • Asks to choose the best Lewis structure for NO₃.

    • Various diagrams presented for evaluation.

Calculating Formal Charge for Nitrogen in NO₃⁻

  • Question: What is the formal charge on N in the nitrite ion (NO₃⁻)?

    • Options:

      • A. -2

      • B. -1

      • C. 0

      • D. +1

      • E. +2

  • Example Calculation:

    • 5−4=+15 - 4 = +1

Exceptions to Lewis Structures

  1. Free radicals: Molecules that have an odd number of electrons.

    • Tend to be highly reactive due to unpaired electrons.

  2. Incomplete octets:

    • Atoms such as boron (B) and beryllium (Be) may form incomplete octets (six electrons).

  3. Expanded octets:

    • Elements in the third row and beyond have the capacity to accommodate up to 12 electrons (and sometimes even 14) due to the availability of d orbitals.

    • Example: Nitrogen monoxide (NO) has 11 total electrons.

    • Correct structure:

      • N+O=5+6=11extelectronsN + O = 5 + 6 = 11 ext{ electrons}

    • Observations:

      • A formal charge of +1 on fluorine (F) is highly unfavorable.

      • Expanding octets generally lowers formal charge.

Best Lewis Structure for PO₄³⁻

  • Considerations for choosing the best construction:

    • Electron Count:

    • 5 (from phosphorus) + 4·6 (from oxygen) + 3 = 32 electrons

    • Octet Rule:

      • Oxygen must complete its octet, phosphorus is expanded.

    • Evaluate formal charges across possible structures, focusing on minimizing overall charge using the rules discussed.

    • Structure options with various formal charge distributions (not depicted).

  • Ensuring that all values and representations are accurately calculated and represented is central to selecting the best Lewis structures.