Polarity of Molecules – Comprehensive Study Notes

Required Documents

• Posted on the module’s Learning Activities page:

  • Periodic Table (reference for atomic symbols & electronegativity values).

  • Electronegativity Values chart (needed for $|\Delta EN|$ calculations).

  • “Summary: Geometry and Polarity for Molecules and Polyatomic Ions”

  • Condenses information from this and the previous presentation.

  • “Geometry and Polarity – Practice Worksheet”

  • Step-by-step framework for all polarity-determination steps.

  • Mirrors the worksheet used in the Molecular Geometry experiment; useful for extra practice.

Context & Scope

• Unit 5 – Part 5 of CHE 111 (Atoms First) textbook, Section 4.6.

• Created by Dr. Lynn Tracey, last updated 15 Aug 2023.

• Goal of this lecture: learn to classify molecules (NOT polyatomic ions) as polar or nonpolar.

  • Polyatomic ions possess an overall formal charge that dominates their behaviour, so polarity analysis is unnecessary for them.

• Importance: Molecular polarity governs many physical/chemical properties (e.g., solubility, boiling point, intermolecular forces). Details explored in later units.

Fundamental Definitions

• Polar molecule

  • Possesses an overall dipole moment: one end is partially positive (δ+) and the other partially negative (δ-).

• Nonpolar molecule

  • Exhibits no overall dipole moment; charge distribution is symmetric.

• Electronegativity (EN)

  • A calculated, dimensionless value indicating how strongly an atom attracts shared electrons in a covalent bond.

  • Difference in EN between two atoms is expressed as $|\Delta EN|$ (absolute value).

  • Larger $|\Delta EN|$ → more polar bond (greater electron-pulling disparity).

  • Course convention:

    • Polar covalent bond if $|\Delta EN| \ge 0.5$.

    • Nonpolar covalent bond if |\Delta EN| < 0.5.

Bond Dipoles

• Every polar bond behaves as a tiny dipole:

  • Negative end: atom with higher EN (electron density ↑).

  • Positive end: atom with lower EN (electron density ↓).

• Two notation styles:

  • Partial-charge symbols: $\delta^+\text{X}—\text{Y}\delta^-$.

  • Arrow/plus sign: $\underset{+}{\text{X}}\longrightarrow \text{Y}$ (arrow points toward δ- end; plus sign marks δ+).

Criteria for Molecular Polarity

• Nonpolar molecule if EITHER condition is met:

  1. No bond dipoles exist (all bonds nonpolar).

  2. Bond dipoles are arranged so they completely cancel.

• Polar molecule if BOTH conditions are met:

  1. At least one polar bond exists (≥ 1 bond dipole).

  2. Bond dipoles do NOT cancel.

Two-Step Decision Process

1. Bond Polarity Check

   • Identify each unique bond type.

   • Compute $|\Delta EN|$ using EN chart.

   • Classify each bond as polar or nonpolar.

2. Dipole-Cancellation Check

   • Examine 3-D molecular geometry (cannot rely on flat Lewis structure).

   • Determine whether the vector sum of all bond dipoles is zero.

   • For simple molecules this is qualitative; rigorous vector math is unnecessary in CHE 111.

Special Rules for Molecules with ONE Central Atom

• Dipoles cancel ONLY IF both of the following are true:

  1. Every electron domain around the central atom is a bonding domain (no lone pairs).

  2. All atoms directly bonded to the central atom are identical (i.e., equal EN values).

• Visual key (used in slides):

  • Black circle = central atom.

  • Blue circles = one set of identical surrounding atoms.

  • Green circles = another set of identical surrounding atoms (if present) with a different EN from the blue set.

  • Small red lobe = non-bonding (lone) pair.

Only Three “Perfect-Cancellation” Geometries (in this course)

• Linear (AX₂)

  • $180^\circ$ bond angle.

  • Two equal-magnitude dipoles in opposite directions → perfect cancellation.

• Trigonal Planar (AX₃)

  • $120^\circ$ bond angles in one plane.

  • Three identical dipoles symmetrically oriented outward/inward → three-way cancellation.

• Tetrahedral (AX₄)

  • $109.5^\circ$ bond angles in 3-D.

  • Four identical dipoles symmetrically oriented → four-way cancellation.

• All other combinations (lone pairs present, different surrounding atoms, or different geometries) result in polar molecules because the dipoles do not fully cancel.

Molecules with MULTIPLE Central Atoms

• Can be nonpolar only if every individual bond dipole in the entire molecule cancels collectively.

• No simple “shortcut rule”; entire 3-D frame must be analysed.

• Reminder: If the molecule contains zero polar bonds, it is automatically nonpolar regardless of size or shape.

Worked Examples (from slides)

Example 1: $\text{AsCl}_3$

• Geometry: Trigonal pyramidal (1 lone pair on As, 3 bonding pairs).

• Bonds present: As–Cl.

  • $|\Delta EN| = 3.0 - 2.0 = 1.0 \Rightarrow$ polar bonds.

• Dipole cancellation?

  • One central atom, 3 atoms + 1 lone pair → cancellation fails.

• Conclusion: POLAR molecule.

• Explanation: Dipoles from the three identical polar bonds are tilted by the lone pair; net vector ≠ 0 → overall dipole present.

Example 2: $\text{CO}_2$

• Geometry: Linear O=C=O.

• Bonds present: C–O.

  • $|\Delta EN| = 3.5 - 2.5 = 1.0 \Rightarrow$ polar bonds.

• Dipole cancellation?

  • Central atom with 2 identical surrounding atoms, 0 lone pairs → criteria for linear cancellation satisfied.

• Conclusion: NONPOLAR molecule.

• Explanation: Two equal polar bonds oriented $180^\circ$ apart cancel exactly → net dipole = 0.

Example 3: $\text{COH}_2$ (Formaldehyde, CH₂O)

• Geometry: Trigonal planar around C (no lone pairs).

• Bonds present & polarity:

  • C–O: $|\Delta EN| = 1.0$ → polar.

  • C–H: $|\Delta EN| = 0.4$ → nonpolar.

• Net dipoles:

  • Only one bond dipole exists (C–O).

  • A single vector cannot cancel itself.

• Conclusion: POLAR molecule.

• Explanation: Unequal EN of attached atoms means symmetry is broken; lone dipole remains unopposed → overall dipole.

Example 4: $\text{SO}_2$

• Geometry: Bent (1 lone pair on S, 2 bonding pairs).

• Bonds present: S–O.

  • $|\Delta EN| = 1.0$ → polar bonds.

• Dipole cancellation?

  • Two polar bonds + lone pair create an angular geometry; vectors not directly opposed.

• Conclusion: POLAR molecule.

• Explanation: Lone pair forces bent shape, preventing linear arrangement; dipoles add to give net moment.

General Summary / Decision Algorithm

1. Identify all bond types.

   • Compute $|\Delta EN|$ for each.

   • Label each bond as polar or nonpolar.

2. If no polar bonds → molecule is nonpolar (done).

3. If ≥ 1 polar bond → inspect molecular geometry.

   • For single-central-atom species, use the three “perfect-cancellation” geometries as a quick test.

   • For others (lone pairs, mixed atoms, multi-central-atom frameworks), draw/visualise the 3-D structure and qualitatively sum the dipole vectors.

4. Conclusions:

   • Polar molecule → at least one polar bond AND incomplete dipole cancellation.

   • Nonpolar molecule → either no polar bonds OR complete cancellation of all dipoles.

Ethical, Philosophical & Practical Implications (highlighted implicitly)

• Understanding polarity underpins responsible handling of chemicals (predicting solubility, toxicity, environmental persistence).

• Encourages molecular-level thinking: appreciating that invisible vector sums dictate macroscopic behaviour.

Numerical / Statistical References & Formula Recap

• Electronegativity cut-offs:

  • $|\Delta EN| \ge 0.5$ → polar covalent.

  • |\Delta EN| < 0.5 → nonpolar covalent.

• Perfect-cancellation bond angles (ideal VSEPR values):

  • Linear: $180^\circ$.

  • Trigonal planar: $120^\circ$.

  • Tetrahedral: $109.5^\circ$.

• Dipole vector representation: arrow points toward δ-, plus sign marks δ+.

Connection to Previous & Upcoming Content

• Builds directly on VSEPR & Molecular Geometry (Unit 5 – Part 4) where shapes/lone pairs were determined.

• Sets the stage for Unit 5 – Part 6 (Sigma & Pi Bonds; Hybrid Orbitals) by emphasising spatial orientation of bonds and electron density.

Practice & Next Steps

• Use the “Geometry and Polarity – Practice Worksheet” to reinforce the two-step method.

• Additional molecules (beyond examples) should be tested to master quick identification.

• Upcoming lecture will explore orbital overlap (σ and π bonds) and hybridisation, further explaining why certain geometries arise and how they affect polarity.

Last slide reminder: “THE END” – proceed to Unit 5 Part 6.