Intermolecular Forces: Dipoles & H-Bonds - W1 AV1

Overview

Topic: Intermolecular forces, focusing on dipole-dipole interactions and hydrogen bonding.
Goals: Define interactions, give examples, and explain effects on boiling point and water solubility.
Emphasis: Interactions occur between molecules, not within a molecule.

Occur between polar molecules that possess permanent dipole moments.
Opposite partial charges on different molecules attract each other.
Example: Acetone — oxygen is partially negative, carbon (adjacent) is partially positive.
Example: Carbon monoxide (CO) — carbon partial positive, oxygen partial negative; two CO molecules attract via dipole-dipole force.
Key point: Dipole-dipole forces act between separate molecules (intermolecular).

A special, stronger type of dipole-dipole interaction.
Occurs when hydrogen is bonded to N, O, or F.
Example: Water — O is partially negative, H is partially positive; O (of one molecule) attracts H (of another) forming an H-bond.
Hydrogen bonds are intermolecular and link separate molecules.

Hydrogen bonding increases boiling point and water solubility.
Examples:
- Ammonia and methanol are highly soluble in water due to H-bonding.
- Methanol mixes completely with water and has a higher boiling point than comparable molecules lacking H-bonds.

Ethanol vs Dimethyl Ether:
- Ethanol has O–H and can hydrogen bond; dimethyl ether cannot (no H on O).
- Ethanol has a much higher boiling point (≈78 °C) than dimethyl ether (≈ -23 °C).
- Ethanol is more soluble in water due to H-bonding.
Ethanol vs 1-Butanol:
- Both have O–H (can hydrogen bond), but 1-butanol has a longer nonpolar hydrocarbon chain.
- 1-Butanol has higher boiling point due to increased London dispersion forces from a larger carbon chain.
- Ethanol is more water-soluble because its nonpolar region is smaller.
Trend: Increasing nonpolar (C–H / C–C) regions decreases water solubility but increases boiling point.

Methanol, ethanol, 1-octanol:
- All can hydrogen bond.
- 1-Octanol has large nonpolar region → not soluble in water, mixes with nonpolar solvents.
- Methanol has highest water solubility (smallest nonpolar region).
- 1-Octanol has the highest boiling point of the three.
Alcohol solubility: small-chain alcohols (methanol, ethanol) are highly soluble; adding CH2 groups reduces solubility.

Constitutional isomers with same formula can have different boiling points.
Pentane vs Neopentane (C5H12):
- Pentane (straight chain) has higher boiling point than neopentane (branched).
- Straight-chain alkanes have greater surface area → stronger London dispersion forces.
General rule: More surface area and less branching → higher London dispersion forces → higher boiling point.

"Like dissolves like": polar solvents dissolve polar solutes; nonpolar solvents dissolve nonpolar solutes.
Molecules composed only of C and H (alkanes) are nonpolar and not soluble in water.
Presence of polar functional groups (e.g., O–H) increases solubility in water.

Dipole-Dipole Interaction: Attraction between oppositely charged ends of polar molecules.
Hydrogen Bond: Strong dipole-dipole interaction when H is bonded to N, O, or F.
London Dispersion Forces: Weak, temporary induced dipole attractions; increase with molecular size and surface area.
Constitutional Isomers: Same molecular formula but different connectivity/structure.
Polar vs Nonpolar: Polar molecules have permanent dipole moments; nonpolar molecules do not.

Identify functional groups (O–H, N–H, C–H) to predict H-bonding and polarity.
Compare similar molecules by counting polar groups and nonpolar carbon chains to estimate solubility and boiling point.
For isomers, evaluate branching and surface area to predict relative boiling points.
Remember: hydrogen bonding raises boiling point and water solubility; larger nonpolar regions lower water solubility but raise boiling point.