CHE125 Study Notes: Lewis Structures and Molecular Geometry 2/19

Overview of Lewis Structures and Charges

  • Topics Covered

    • Electronic Configuration

    • Ion Formula

    • Lewis Structures

    • Formal Charge Calculation

  • Laboratory Sessions

    • No labs this week due to a holiday.

    • Recitation sessions are ongoing.

Lewis Structures

  • Introduction

    • Lewis Structures are visual representations of molecular structures that show how atoms are bonded and the arrangement of electrons.

    • Considerations include valence electrons, bond arrangements, and the formal charges of atoms.

  • Examples

    • NCO (negative one charge)

    • PO₄ (negative three charge)

    • Total valence electrons for PO₄: 29 (5 from phosphorus + 24 from four oxygens, adjusted for three extra electrons due to the overall -3 charge)

  • Drawing Lewis Structures

    • Start with total electron count:

    • For NCO: 16 electrons (1 from nitrogen, 6 from carbon, and 6 from each oxygen)

    • Bond types are added (single or double as needed).

Steps to Create Lewis Structures

  1. Count Total Valence Electrons

    • For NCO:

      • N (5) + C (4) + O (6) = 15; with one bond, total becomes 16 after accounting for the charge.

  2. Distribute Electrons

    • Start by forming single bonds between the central atom and surrounding atoms.

    • Complete octets for outer atoms (two electrons per bond).

    • Use remaining electrons to form double bonds if octets are not satisfied.

  3. Formal Charge Calculation

    • Definition: Formal charge = (Valence electrons of the atom) - (Nonbonding electrons) - 0.5*(Bonding electrons)

    • Apply formal charge calculations for each atom within the structure to verify overall charge consistency.

Formal Charge Examples

  • Calculation for Different Atoms in Lewis Structures

    • Oxygen:

    • Normal electrons: 6, Electrons after bonding: 6 (if involved in double or single bonds), resulting in a formal charge of 0,

    • Or: Formal charge = 6 - 7 = -1 if it is holding one extra paired bond.

    • Nitrogen:

    • Normal = 5, involved in bonding = 6

      • Example result: 5 - 6 = -1.

    • Carbon: Always adjusts based on shared electron counts from double or single bonds.

Strategies for Better Lewis Structures

  • Explore alternate bonding configurations if formal charges yield unfavorable results:

    • Shared Bonds:

    • Shift electron pairs to create double/triple bonds as needed to obtain a more favorable charge setup.

    • Identify major vs. minor resonance structures:

    • Structures exist in equilibrium, with one usually being the dominant (most stable) form.

    • Resonance: A term used to describe the phenomenon where two or more possible structures describe a single molecule, with one shape being more stable than the others.

      • e.g., comparing structures A, B, and C based on formal charge distributions to determine which is the most stable based on the least charge discrepancy.

Electronegativity and Polarity

  • Concept of Electronegativity

    • Definition: Atoms’ tendency to pull electrons toward themselves when forming bonds.

    • Fluorine is the most electronegative element, with stronger pulls compared to others.

    • Polarity

    • Polar Bonds: Arise from differences in electronegativity, leading to uneven electron sharing between atoms.

    • To represent polarity in bonds, arrows are used pointing towards the more electronegative atom, indicating a partial negative charge at that end.

    • Identifying Polarization

    • Assess bonds within structures:

      • Compare differences in electronegativity between bonded atoms.

      • Example: H-O is polar due to O's greater electronegativity, while C-H displays minimal polarity due to close electronegativities.

Molecular Geometry and VSEPR Theory

  • VSEPR Theory: Valence Shell Electron Pair Repulsion theory explains the three-dimensional shape adopted by molecules based on electron interactions:

    • Straightforward geometric arrangements are determined by the number of electron pair domains around a central atom.

    • Geometries include:

    • Linear (2 domains)

    • Trigonal planar (3 domains)

    • Tetrahedral (4 domains)

  • Molecular Geometry: Evaluates shapes excluding lone pairs on the central atom, which defines the visible structure of a molecule:

    • E.g., NH₃ would exhibit a pyramidal shape due to a lone pair altering its overall geometry.

    • Key Takeaways:

    • Molecular geometry often deviates from counting all electrons due to the invisibility of lone pairs, hence producing varied shapes and properties.

    • Importance of charge, octets, electronegativity, and structural resonance all contribute to determining the most stable molecular configuration of a compound.