Intermolecular Forces Review for AP Chem

What You Need to Know

Intermolecular forces (IMFs) are attractions between particles (molecules/atoms/ions). They don’t change chemical identity, but they strongly control physical properties: boiling point, melting point, vapor pressure, viscosity, surface tension, and solubility.

The core idea (AP Chem-level)

• Stronger IMFs ⟶ particles are harder to separate ⟶ higher boiling point and melting point, lower vapor pressure, higher $\Delta H_{vap}$ (and usually $\Delta H_{fus}$), higher viscosity and surface tension.
• IMFs are about electrostatic attraction: partial charges, permanent dipoles, induced dipoles, and ions.

Critical reminder: Intramolecular forces (covalent/ionic/metallic bonds) are usually much stronger than intermolecular forces. Don’t mix them up when comparing boiling points of similar-sized molecules.

The “big 5” IMFs you need

1. London dispersion forces (LDF) (instantaneous/induced dipoles) — present in all substances.
2. Dipole–dipole — between polar molecules.
3. Hydrogen bonding — a strong subtype of dipole–dipole.
4. Ion–dipole — key for dissolving ionic compounds in polar solvents.
5. (Less common but testable) Dipole-induced dipole / ion-induced dipole.

Step-by-Step Breakdown

Use this whenever you’re asked to identify IMFs, rank boiling points, rank vapor pressures, or predict solubility.

A. Identify the dominant IMFs (molecule vs molecule)

1. Determine the particles present
   • Neutral molecules? Ions? A mixture (solute + solvent)?
2. Check for hydrogen bonding
   • You have H-bond donors if you see $\text{H–N}$, $\text{H–O}$, or $\text{H–F}$.
   • You have H-bond acceptors if you have $\text{N}$, $\text{O}$, or $\text{F}$ with lone pairs (most cases).
   • If donor + acceptor are available between molecules ⟶ hydrogen bonding is significant.
3. If no H-bonding, check polarity (dipole–dipole)
   • Use molecular geometry + bond polarity (net dipole?).
   • Polar molecules ⟶ dipole–dipole + LDF.
   • Nonpolar molecules ⟶ LDF only.
4. Estimate LDF strength (this often decides rankings)
   • Larger molar mass / more electrons ⟶ stronger LDF.
   • Greater surface area (less branching, more “spread out”) ⟶ stronger LDF.
   • More polarizable electron clouds (heavier atoms, $\pi$ systems) ⟶ stronger LDF.

B. Rank physical properties (boiling point, vapor pressure, etc.)

1. Compare IMF types first (H-bonding > dipole–dipole > LDF for similarly sized molecules).
2. If IMF types are similar, use LDF trends:
   • Higher molar mass / more surface area ⟶ higher bp.
3. Connect to the property asked:
   • Stronger IMFs ⟶ higher bp/mp/viscosity/surface tension, lower vapor pressure.

C. Predict solubility / miscibility

1. Determine solute’s dominant interactions (polar? H-bonding? ionic?).
2. Determine solvent’s dominant interactions.
3. Apply “like dissolves like”:
   • Polar + polar (especially H-bonding) ⟶ good solubility.
   • Nonpolar + nonpolar ⟶ good solubility.
   • Ionic + polar ⟶ often soluble via ion–dipole.

Mini worked method (quick)

Rank boiling points: $\text{CH}_4$, $\text{CH}_3\text{OH}$, $\text{CH}_3\text{F}$

• $\text{CH}_4$: nonpolar ⟶ LDF only (weak).
• $\text{CH}_3\text{F}$: polar ⟶ dipole–dipole + LDF.
• $\text{CH}_3\text{OH}$: has $\text{O–H}$ ⟶ hydrogen bonding + others.
  Result: $\text{CH}_3\text{OH} > \text{CH}_3\text{F} > \text{CH}_4$ (bp).

Key Formulas, Rules & Facts

IMF types (what causes them + relative strength)

Strength trend (typical, not absolute)

For similarly sized species:

• Ion–dipole (often very strong in solutions)
• Hydrogen bonding
• Dipole–dipole
• LDF (but can be large for big molecules)

Trap: Huge LDF in large nonpolar molecules can beat dipole–dipole in smaller polar molecules.

LDF / polarizability rules (high yield)

• Polarizability = how easily the electron cloud distorts.
• Increases with:
  • More electrons (down a group)
  • Larger atomic/ionic radius (generally)
  • More extended shape (less branching)
  • More surface area in contact (linear chains > branched isomers)

Property relationships (memorize the direction)

Vapor pressure / temperature link (Clausius–Clapeyron)

Used when comparing liquids or calculating vapor pressure changes:

$\ln\left(\frac{P_2}{P_1}\right) = -\frac{\Delta H_{vap}}{R}\left(\frac{1}{T_2}-\frac{1}{T_1}\right)$

• Stronger IMFs ⟶ larger $\Delta H_{vap}$ ⟶ vapor pressure rises more slowly with temperature.
• Normal boiling point: temperature where vapor pressure $= 1\ \text{atm}$.

Hydrogen bonding specifics (quick rules)

• Donors: molecules containing $\text{H–N}$, $\text{H–O}$, $\text{H–F}$.
• Acceptors: lone-pair-bearing $\text{N}$, $\text{O}$, $\text{F}$ (in most neutral molecules).
• Not hydrogen bonding (common AP trick):
  • $\text{H}$ bonded to $\text{C}$ (like in alcohol-adjacent C–H) is not an H-bond donor.
  • Molecules like $\text{CH}_3\text{OCH}_3$ can **accept** H-bonds but cannot **donate** them (no $\text{O–H}$).

Examples & Applications

Example 1: Rank boiling points (isomers / surface area)

Rank bp: n-pentane, isopentane, neopentane (all $\text{C}_5\text{H}_{12}$).

• All nonpolar ⟶ LDF only.
• LDF strength tracks surface area: linear > branched.
• Result: n-pentane > isopentane > neopentane.

Exam vibe: They love branching vs linear comparisons because polarity is identical.

Example 2: Compare bp when polarity competes with molar mass

Which has higher bp: $\text{H}_2\text{S}$ or $\text{H}_2\text{O}$?

• $\text{H}_2\text{O}$: hydrogen bonding (strong).
• $\text{H}_2\text{S}$: no true hydrogen bonding (S is not N/O/F), has dipole–dipole + LDF.
• Result: $\text{H}_2\text{O}$ is much higher bp due to H-bonding.

Key insight: H-bonding can dominate even when the other molecule is heavier.

Example 3: Vapor pressure ranking at the same temperature

At $25^\circ\text{C}$, rank vapor pressure: ethanol ($\text{CH}_3\text{CH}_2\text{OH}$), dimethyl ether ($\text{CH}_3\text{OCH}_3$), propane ($\text{C}_3\text{H}_8$).

• Ethanol: hydrogen bonding ⟶ strongest IMFs ⟶ lowest vapor pressure.
• Dimethyl ether: polar, no donor H ⟶ dipole–dipole (and can accept H-bonds but not self-donate) ⟶ intermediate.
• Propane: nonpolar, small ⟶ weakest IMFs ⟶ highest vapor pressure.
• Result: propane > dimethyl ether > ethanol (vapor pressure).

Example 4: Solubility (ion–dipole vs “like dissolves like”)

Why does $\text{NaCl}$ dissolve in water but not in hexane?

• Water is polar and can orient its dipoles around $\text{Na}^+$ and $\text{Cl}^-$ ⟶ strong ion–dipole attractions stabilize ions in solution.
• Hexane is nonpolar ⟶ cannot stabilize separated ions effectively.

Exam vibe: Mention ion–dipole explicitly for ionic solids in polar solvents.

Common Mistakes & Traps

1. Mixing up intramolecular vs intermolecular forces

   • What goes wrong: You say “covalent bonds are stronger, so it has higher bp.”
   • Why wrong: Boiling breaks IMFs, not covalent bonds.
   • Fix: For bp/mp/phase, talk only about IMFs and structure.

2. Saying “nonpolar = no forces”

   • What goes wrong: You claim nonpolar molecules have no attractions.
   • Why wrong: LDF are always present.
   • Fix: Always write “LDF only” for nonpolar molecular substances.

3. Calling any molecule with $\text{H}$ a hydrogen-bonder

   • What goes wrong: You treat $\text{CH}_3\text{Cl}$ or $\text{CH}_4$ as H-bonding.
   • Why wrong: Hydrogen bonding needs $\text{H}$ directly bonded to N/O/F.
   • Fix: Look specifically for $\text{O–H}$, $\text{N–H}$, $\text{F–H}$.

4. Forgetting shape when deciding polarity

   • What goes wrong: You see polar bonds and assume the molecule is polar.
   • Why wrong: Dipoles can cancel (ex: $\text{CO}_2$, $\text{CCl}_4$).
   • Fix: Determine geometry; ask “Do bond dipoles cancel?”

5. Ignoring LDF trends when IMF types match

   • What goes wrong: You rank bp of alkanes incorrectly because “all are nonpolar.”
   • Why wrong: LDF strength varies hugely with size and shape.
   • Fix: Compare molar mass and surface area/branching.

6. Assuming dipole–dipole always beats LDF

   • What goes wrong: You predict a small polar molecule must have higher bp than a large nonpolar one.
   • Why wrong: Large nonpolar molecules can have very strong LDF.
   • Fix: If sizes differ a lot, LDF may dominate.

7. Confusing “can H-bond with water” vs “self H-bonding”

   • What goes wrong: You assume ethers have the same bp behavior as alcohols.
   • Why wrong: Ethers can accept H-bonds but cannot donate; pure ether lacks strong H-bond networks.
   • Fix: For bp of a pure substance, ask whether molecules can H-bond to each other (donor + acceptor).

8. Overinterpreting melting point trends

   • What goes wrong: You apply “stronger IMFs ⟶ higher mp” without considering packing.
   • Why wrong: mp depends on IMFs and crystal packing/symmetry.
   • Fix: Use mp trends cautiously; bp trends are usually more straightforward for IMF comparisons.

Memory Aids & Quick Tricks

Quick Review Checklist

• You can name and recognize: LDF, dipole–dipole, hydrogen bonding, ion–dipole, induced-dipole interactions.
• You remember LDF are always present and get stronger with more electrons, bigger size, and more surface area.
• You check molecular shape to decide if a molecule is polar or nonpolar.
• You identify hydrogen bonding only when $\text{H}$ is bonded to N/O/F (and there’s an acceptor available).
• You can rank properties:
  • Stronger IMFs ⟶ bp/mp/viscosity/surface tension increase
  • Stronger IMFs ⟶ vapor pressure/volatility decrease
• You apply “like dissolves like” and call out ion–dipole for ionic solubility in polar solvents.
• You recognize when mp can be “weird” due to packing, but bp is usually the cleaner IMF comparison.

You’ve got this—if you stay systematic (IMF type ⟶ LDF trends ⟶ property direction), these questions become free points.