Polarity of Molecules - Detailed Study Notes

Polarity of Molecules

Introduction

  • Instructor: Ms. Katherine Avante, LPT

Definition of Polarity

  • Polarity refers to the nature of molecules based on the sharing of electrons among the atoms.
  • It can be classified as:
    • Equal sharing of electrons among the atoms of a molecule (Nonpolar).
    • Unequal sharing of electrons (Polar).

Polar Molecule

  • In a polar molecule, there is an unequal or asymmetrical distribution of electrons among the atoms.

Electron Configuration

  • Example: Water molecule structure
    • Representation:
    • H
    • H
    • O
  • Valence Electrons:
    • Definition: An electron in the outer shell associated with an atom, which can participate in the formation of a chemical bond if the outer shell is not closed.
    • For hydrogen, there are 2 valence electrons. For oxygen, there are 6 valence electrons.

Octet Rule

  • States that atoms tend to gain, share, or transfer electrons to attain a stable configuration of 8 valence electrons (octet).

Electronegativity

  • Definition: The relative ability of an atom to attract electrons in a chemical bond toward itself.
  • Analogy: An individual’s strength in a romantic relationship can be compared to an atom's strength in attracting electrons. ("I'm falling for him! He’s so strong!")

Comparison of Electronegativity

  • Oxygen is denser than Hydrogen.
  • Molecules have a dipole (two poles) indicating positive and negative regions.
  • Oxygen is more electronegative than Hydrogen.

Tug of War Analogy

  • Visualizes electron sharing as a tug of war, where the shared electron pairs are attracted more strongly toward the more electronegative atom, resulting in unequal sharing.

Nonpolar Molecule

  • In a nonpolar molecule, there is equal or symmetrical distribution of electrons among the atoms.
  • Example: Oxygen gas (O₂)
    • Representation: O=O
    • Exhibits equal ability to attract.
    • No partial charge or dipole movement present.

Determining Polarity Based on Electronegativity

  • Types of Bonds based on Electronegativity:
    • Pure (nonpolar) covalent bond: Electrons shared equally
    • Polar covalent bond: Electrons shared unequally
    • Ionic bond: Electron transferred

Electronegativity Values

  • Table of electronegativity values and example differences:
    • H: 2.20
    • Li: 0.98
    • Na: 0.93
    • Cl: 3.0
    • Difference for Ionic bond: >2.1 (e.g., sodium chloride NaCl, with a difference of 2.1)

Polar vs. Nonpolar Covalent Bond

Polar Covalent Bond
  • Features:
    • Unpaired valence electrons
    • Unequal sharing of valence electrons
Nonpolar Covalent Bond
  • Features:
    • Unpaired valence electrons
    • Equal sharing of valence electrons
Ionic Bond
  • Features:
    • Transfer of electrons between atoms

Periodic Table and Electronegativity

  • Electronegativity increases across the periodic table due to atomic structure considerations.
  • Example values across groups and periods.

Electrostatic Attraction

  • Definition: Phenomena where opposite charges are pulled toward one another.
  • The negative charge of electrons is attracted to the positive charge presented in the nucleus.

Electronegativity Difference and Bond Type

  • Ranges of electronegativity differences determine bond types:
    • Nonpolar covalent bond: 0-0.4
    • Polar covalent bond: 0.5-2.0
    • Ionic bond: >2.1

Molecular Geometry

Definition
  • Relates to the three-dimensional arrangement of atoms in a molecule.
Valence Shell Electron Pair Repulsion (VSEPR) Theory
  • Basic concepts to remember:
    • Lewis Electron Dot Structure (LEDS) is used to predict 3-D molecular geometry based on the number of valence shell electron bond pairs in a molecule or ion.

Steps for Drawing Lewis Structures

  1. Count all the valence electrons.
  2. Determine the central atom (usually the element present in the least number).
  3. Draw single bonds to the central atom.
  4. Place remaining electrons on the other atoms making sure each achieves an octet or duet as required.
  5. Convert lone pairs into double or triple bonds where applicable.

VSEPR Theory Key Ideas

  1. Electron pairs stay as far apart from each other as possible to minimize repulsions.
  2. The shapes are determined by the number of bond pairs and lone pairs around the central atom.
  3. Multiple bonds are treated as single bonds for predictions of shape.
  4. Lone pairs occupy more space than bond pairs.

Determining Molecular Shape Examples

Linear
  • 2 bond pairs around a central atom, 180° bond angle
Trigonal Planar
  • 3 bond pairs, 120° bond angle
Bent
  • 2 bond pairs and 1 lone pair around the atom, <120° bond angle for 3 total regions.
Tetrahedral
  • 4 bond pairs around a central atom, 109.5° bond angle
Trigonal Pyramidal
  • 3 bond pairs and 1 lone pair, <109.5° bond angle
Trigonal Bipyramidal
  • 5 bond pairs, specific angles dependent on axial and equatorial positions
Octahedral
  • 6 bond pairs, 90° bond angles

Polar vs. Nonpolar Molecules

  • A nonpolar molecule is symmetrical with equal sharing of electrons and no dipole moment.
  • A polar molecule is asymmetrical with unequal sharing of electrons and has a defined dipole moment.

Summary of Key Concepts

  1. Polar Bond: Refers to the manner in which electrons are shared between atoms.
    • If the bond is polar, the molecule may be either polar or nonpolar.
    • If the bond is nonpolar, the molecule is automatically considered nonpolar.
  2. Checking Polarity: Determine the shape of the molecule and account for dipole moment vectors to ascertain if they cancel out.
  3. Conclusion: Discussion of how molecular geometry and polarity influence biological, chemical, and physical properties of substances.