Bond Polarity & Electronegativity Notes

Bond Polarity & Electronegativity

  • Electronegativity (EN)
    • A dimensionless measure of an atom’s ability to attract shared electrons.
    • The higher the EN, the stronger the pull on bonding electrons.
  • Bond polarity depends on the EN difference \Delta EN = |ENA - ENB|.
    • \Delta EN = 0 → perfectly even sharing.
    • Larger \Delta EN → greater electron displacement, producing partial charges \delta^+ and \delta^-.
  • Key rule ("bigger gap, bigger pull"):
    • “The greater the difference in electronegativity, the more polar the bond.”

Electronegativity‐Difference Classification

  • (Typical cutoffs; refer to your course’s exact table.)
    • 0 \le \Delta EN \le 0.4 → Non-polar covalent (electron cloud evenly distributed).
    • 0.4 < \Delta EN \le 1.7 → Polar covalent (partial charges form, but electrons still shared).
    • \Delta EN > 1.7 → Ionic (electron transfer, not sharing).

Worked Example

  1. Look up chlorine’s EN: EN_{Cl} = 3.0.
  2. For a Cl–Cl bond (diatomic Cl_2): \Delta EN = |3.0 - 3.0| = 0.
  3. Classify: \Delta EN = 0 → non-polar covalent.
    • Electrons are equally distributed; there is no net positive or negative end.

Visual / Metaphor Mentioned

  • “Polar bear stranded on an iceberg, dissolving in water”
    • Used to dramatize polarity (anything polar tends to interact with polar water).
  • “Regular bear” contrasts with “polar bear” to reinforce the idea of non-polar vs. polar species.

Molecular Polarity vs. Bond Polarity

  • A molecule can contain polar bonds yet be non-polar overall if its shape is symmetrical (vector sum of dipoles cancels).
    • Determined using VSEPR (Valence Shell Electron Pair Repulsion) theory (spoken in transcript as “V-expert”).
  • Example idea alluded to (not explicitly named in clip): CO_2 has two polar C=O bonds, but linear geometry → net dipole =0.

Consequences & Physical Properties

  • Net polarity influences intermolecular forces:
    • Stronger dipole–dipole or hydrogen bonding → higher boiling/melting points.
    • Non-polar molecules rely on weaker London dispersion forces.
  • Thus, bond-polarity analysis helps predict bp/mp trends, solubility ("like dissolves like"), and other physical behavior.

Key Takeaways

  • Always start by calculating \Delta EN.
  • Consult the cut-off table to label the bond (non-polar, polar covalent, ionic).
  • Use VSEPR shapes to decide if individual bond dipoles cancel.
  • Remember: Polarity is a vector property—both magnitude (from \Delta EN) and direction (from geometry) matter.
  • Physical properties (bp, mp, solubility) correlate strongly with overall polarity.