Microscopic World 2 Revision Notes (Section A & Section B)

Section A

• Multiple Choice (MC) Answers
  • 1–5: AACDD
  • 6–10: BDDAD
  • 11–15: DBADD
  • 16–17: BD

Section B

LQ 1

• Concept: Hydrogen bonding should be shown as a bond formed between a hydrogen attached to an electronegative atom and a lone pair on another electronegative atom.
  • The illustration should include: the hydrogen bond; the lone pair on the acceptor atom (e.g., oxygen) donating to the hydrogen; and the resulting partial charges (δ+ on H, δ− on the acceptor electronegative atom).
• Relevance: Hydrogen bonding is a key intermolecular force that explains higher boiling points for molecules with O–H, N–H, or F–H groups compared with similar non-hydrogen-bonded analogues.

LQ 2

• (a) Molecular geometries around a central atom to identify:
  • Tetrahedral
  • Trigonal pyramidal
  • V‑shape / bent
• (b) C–H bond polarity:
  • The C–H bond is polar because there is an electronegativity difference between carbon and hydrogen.
  • On the Pauling scale, C is slightly more electronegative than H, leading to a small dipole with δ− on C and δ+ on H.
  • Net molecular polarity depends on geometry and the sum of dipoles; methane (CH₄) is non-polar due to its tetrahedral symmetry and cancellation of C–H dipoles.

LQ 3

• (a) CO₂
  • C=O bonds are polar because oxygen is more electronegative than carbon.
  • CO₂ is a linear, symmetric molecule; the polar bonds cancel each other, giving a net dipole moment of zero.
  • Therefore, CO₂ is non-polar (despite having polar bonds).
  • Formula reference: $ext{CO}_2$
• (b) SO₂
  • S=O bonds are polar because oxygen is more electronegative than sulfur.
  • SO₂ has a V‑shaped (bent) geometry; the bond dipoles do not cancel out.
  • Result: SO₂ is polar.
  • Note from transcript: it states "S=O is polar bond since S is more electronegative than C" (likely a typographical error in the source; the correct comparison is O vs S). The SO₂ polarity arises from O's higher electronegativity relative to S in the S–O bonds, combined with the bent shape.

LQ 4

• Boiling point (bp) trends for the listed compounds:
  • Butane < Propanone < Propan-2-ol < Propan-1-ol
  • For bp comparison notes:
  • Butane vs Propanone:
    • Butane is non-polar; van der Waals (London dispersion) forces are weaker between butane molecules than the dipole–dipole interactions in propanone, so butane has the lower bp.
    • Formula references: butane ≈ $ext{C}<em>4 ext{H}</em>{10}$; propanone ≈ $ext{CH}<em>3 ext{COCH}</em>3$ or $ext{C}<em>3 ext{H}</em>6 ext{O}$.
  • Propanone vs Propan-2-ol:
    • Propan-2-ol has hydrogen bonding (–OH group), yielding stronger intermolecular forces than the van der Waals and dipole forces in propanone.
    • Therefore, bp(propan-2-ol) > bp(propanone).
    • Note: In the transcript, the order is propanone < propan-2-ol, which aligns with this explanation.
  • Propan-2-ol vs Propan-1-ol:
    • Propan-2-ol isomer has hydrogen bonding but a smaller surface area compared to propan-1-ol; propan-1-ol generally has a higher bp due to a larger molecular surface area and extended van der Waals forces and hydrogen-bonding network along its chain.
    • Therefore, bp(propan-2-ol) < bp(propan-1-ol).
• Summary of the reasoning the transcript provides:
  • bp hierarchy reflects the balance of intermolecular forces: London dispersion (non-polar) vs dipole–dipole (polar) vs hydrogen bonding (strongest among these for O–H-containing alcohols).
  • The transcript explicitly states: "Butane < propanone < propan-2-ol < propan-1-ol" and then explains the comparison using the type and strength of intermolecular forces for each pairing.

Notes and clarifications

• Throughout, the transcript emphasizes polarity and molecular geometry as key determinants of physical properties such as boiling point and polarity.
• Where the transcript contains a potential inconsistency (S vs C in LQ 3(b) note), the correct chemical reasoning is included above, but the original wording is preserved in the notes with a corrective comment.
• Formulas referenced in the notes include:
  • Methane: $ext{CH}_4$
  • Butane: $ext{C}<em>4 ext{H}</em>{10}$
  • Propanone (acetone): $ext{CH}<em>3 ext{COCH}</em>3$ or $ext{C}<em>3 ext{H}</em>6 ext{O}$
  • CO₂: $ext{CO}_2$
  • SO₂: $ext{SO}_2$
• Key concepts linked to foundational ideas:
  • Polarity and dipole moments, effect of molecular symmetry, and how geometry affects overall polarity.
  • Types of intermolecular forces: London dispersion, dipole–dipole, and hydrogen bonding, and how these influence boiling points.
  • The role of electronegativity differences in determining bond polarity (e.g., C–H vs C=O vs S–O).