Why IMFs Matter

Physical properties affected by IMF strength:

1. Boiling point: Stronger IMF → higher BP
2. Melting point: Stronger IMF → higher MP
3. Vapor pressure: Stronger IMF → lower VP
4. Viscosity: Stronger IMF → higher viscosity
5. Surface tension: Stronger IMF → higher surface tension
6. Solubility: "Like dissolves like" (similar IMFs)

Three Main Types of IMFs

1. London Dispersion Forces (LDF)

Also called: Van der Waals forces, induced dipole-induced dipole

Definition: Temporary attractive forces caused by temporary dipoles in atoms/molecules

Mechanism:

1. Electrons in atom/molecule move randomly
2. At any instant, electron distribution may be asymmetric
3. Creates temporary (instantaneous) dipole
4. This dipole induces dipole in neighboring molecule
5. Temporary dipoles attract each other

Key points:

• Present in ALL molecules (only IMF in nonpolar molecules)
• Weakest type of IMF
• Very short-lived (constantly forming and breaking)

Strength factors:

a) Molar mass / Number of electrons:

• More electrons → stronger LDF
• Larger electron cloud → more polarizable
• Bigger molecules → stronger LDF

Example: I₂ > Br₂ > Cl₂ > F₂ (increasing LDF strength)

b) Molecular shape:

• More surface area contact → stronger LDF
• Linear molecules > compact spherical molecules

Example: n-pentane (linear) vs. neopentane (spherical)

• Same formula (C₅H₁₂), but n-pentane has higher BP due to more surface contact

Strength range: 0.1 - 2 kJ/mol (typically)

Examples of molecules with ONLY LDF:

• Noble gases (He, Ne, Ar, Kr, Xe)
• Nonpolar molecules (H₂, N₂, O₂, CH₄, CO₂, CCl₄)

2. Dipole-Dipole Forces

Definition: Attractive forces between permanent dipoles in polar molecules

Mechanism:

1. Polar molecules have permanent dipoles (δ+ and δ-)
2. Positive end of one molecule attracts negative end of another
3. Molecules align: δ+ ↔ δ-

Requirements:

• Molecule must be polar (permanent dipole moment)
• Asymmetric molecular geometry
• Polar bonds that don't cancel

Key points:

• Only in polar molecules
• Stronger than London dispersion forces
• Molecules with dipole-dipole also have LDF!

Strength factors:

• Greater polarity (larger dipole moment) → stronger force
• Electronegativity difference matters

Strength range: 5 - 25 kJ/mol (typically)

Examples:

• HCl, HBr, HI (polar molecules)
• SO₂ (bent, polar)
• CH₃Cl (polar)
• Acetone (CH₃COCH₃)

Effect on properties:

• Polar molecules have higher BP than nonpolar molecules of similar mass
• Example: CH₃Cl (BP = -24°C) vs. CH₄ (BP = -161°C)

3. Hydrogen Bonding

Definition: Special, extra-strong dipole-dipole interaction when H is bonded to N, O, or F

Requirements (both must be met):

1. H bonded directly to N, O, or F (highly electronegative atoms)
2. Another molecule with lone pair on N, O, or F

Mechanism:

• H bonded to N/O/F becomes very δ+
• Small H atom allows close approach to lone pair on another N/O/F
• Strong electrostatic attraction
• Partially covalent character

Representation:

$R-O-H ... O-R$

The dotted line (...) represents hydrogen bond

Key points:

• Strongest type of intermolecular force
• Much stronger than regular dipole-dipole
• Responsible for water's unique properties
• Directional (involves specific geometry)

Strength range: 10 - 40 kJ/mol (strongest IMF)

Common examples:

• Water (H₂O): Each molecule can form 4 H-bonds
  • 2 H atoms (donors)
  • 2 lone pairs on O (acceptors)
• Ammonia (NH₃): Can form 4 H-bonds
  • 3 H atoms (donors)
  • 1 lone pair on N (acceptor)
• Hydrogen fluoride (HF): Extensive H-bonding
• Alcohols (R-OH): One H-bond donor, one acceptor
• Carboxylic acids (R-COOH): Strong H-bonding (often form dimers)

Special note about water:

• Ice has open structure due to H-bonding
• Each H₂O forms 4 H-bonds in tetrahedral arrangement
• Ice is less dense than liquid water (unusual!)
• Maximum density at 4°C

Effect on properties:

• Dramatically higher BP than expected
• Example: H₂O (BP = 100°C) vs. H₂S (BP = -60°C)
• Similar mass, but H₂O has H-bonding, H₂S doesn't

Comparing IMF Strengths

General strength order:

$\text{London Dispersion} < \text{Dipole-Dipole} < \text{Hydrogen Bonding}$

However:

• Very large nonpolar molecules (strong LDF) can have higher BP than small polar molecules
• Example: C₁₀H₂₂ (nonpolar, high BP) vs. CH₃OH (polar, lower BP)

For similar molecules:

$\text{Hydrogen bonding} > \text{Dipole-dipole} > \text{London dispersion}$

Identifying IMFs - Step by Step

Step 1: Determine if molecule is polar or nonpolar

• Draw Lewis structure
• Determine geometry (VSEPR)
• Check for symmetry

Step 2: Identify IMFs present

If nonpolar:

• Only London dispersion forces
• Examples: CH₄, CCl₄, CO₂, noble gases

If polar (but no N-H, O-H, or F-H bonds):

• Dipole-dipole + London dispersion
• Examples: HCl, CH₃Cl, SO₂

If has N-H, O-H, or F-H bonds:

• Hydrogen bonding + dipole-dipole + London dispersion
• Examples: H₂O, NH₃, CH₃OH, HF

Step 3: Identify strongest IMF

• This determines physical properties

Effect on Boiling Points

Boiling point trends:

Trend 1: Within same type of IMF

Larger molar mass → Higher BP (stronger LDF)

Example - Halogens (all nonpolar, only LDF):

• F₂ (38 g/mol): BP = -188°C
• Cl₂ (71 g/mol): BP = -34°C
• Br₂ (160 g/mol): BP = 59°C
• I₂ (254 g/mol): BP = 184°C

Trend 2: Different types of IMF

H-bonding > Dipole-dipole > LDF only

Example - Similar molar masses:

• CH₄ (16 g/mol, LDF only): BP = -161°C
• NH₃ (17 g/mol, H-bonding): BP = -33°C
• H₂O (18 g/mol, H-bonding): BP = 100°C

Trend 3: Period trends

Hydrides of Group 14-17:

Group 14 (no polarity): CH₄ < SiH₄ < GeH₄ < SnH₄ (normal trend)

Group 15-17 (H-bonding possible):

• NH₃ > PH₃ < AsH₃ < SbH₃ (anomaly due to H-bonding)
• H₂O > H₂S < H₂Se < H₂Te (anomaly due to H-bonding)
• HF > HCl < HBr < HI (anomaly due to H-bonding)

Explanation: Period 2 hydrides (NH₃, H₂O, HF) have H-bonding, dramatically increasing BP

Molecular Comparisons

Example 1: CH₄ vs. NH₃ vs. H₂O

Trend: H₂O > NH₃ > CH₄

Why: H₂O has strongest H-bonding (2 donors, 2 acceptors)

Example 2: n-Pentane vs. Neopentane

Same formula (C₅H₁₂), different shapes:

n-Pentane: Linear chain

• More surface area contact
• Stronger LDF
• BP = 36°C

Neopentane: Compact, spherical

• Less surface area contact
• Weaker LDF
• BP = 10°C

Trend: Linear > branched for same molecular formula

Special Cases

1. Intramolecular H-bonding

Example: Salicylic acid

• H-bond within same molecule
• Lower BP than expected
• Less intermolecular H-bonding available

2. Network Covalent Solids

Examples: Diamond, SiO₂, SiC

• Covalent bonds throughout (NOT IMFs)
• Extremely high melting points
• Very hard

3. Metallic Bonding

Metals: Strong bonding, but not IMF

• Delocalized electrons
• High melting points (generally)

Summary Table

Note: Molecules with dipole-dipole or H-bonding also have LDF!

Applications

1. Predicting physical states: Stronger IMF → likely solid/liquid at room temperature
2. Boiling/melting points: Stronger IMF → higher BP/MP
3. Solubility: Similar IMFs → soluble ("like dissolves like")
4. Biological systems: H-bonding in DNA, proteins, cell membranes
5. Materials science: IMFs affect polymer properties, adhesion, coatings

IMFs present:

• London Dispersion Forces (LDF) only
• All molecules have LDF
• No permanent dipole → no dipole-dipole
• No O-H, N-H, or F-H → no hydrogen bonding

Answer for (a): London dispersion forces only

Step 2: Analyze HCl (hydrogen chloride)

Lewis structure: H-Cl with 3 lone pairs on Cl

Molecular geometry: Linear (diatomic)

Polarity check:

• H-Cl bond is polar (ΔEN = |3.0 - 2.1| = 0.9)
• Only one bond (diatomic)
• Polar molecule with permanent dipole
• δ+ on H, δ- on Cl

IMFs present:

1. London Dispersion Forces (LDF) ✓

   • All molecules have LDF

2. Dipole-Dipole Forces ✓

   • HCl is polar
   • Permanent dipole moment
   • Molecules align: H(δ+) ↔ Cl(δ-)

3. Hydrogen Bonding? ✗

   • Has H, but H is bonded to Cl (not N, O, or F)
   • Cl is not electronegative enough
   • No hydrogen bonding

Answer for (b): Dipole-dipole forces and London dispersion forces

Step 3: Analyze NH₃ (ammonia)

Lewis structure: N with 3 H atoms and 1 lone pair

Molecular geometry: Trigonal pyramidal (SN = 4, 1 lone pair)

Polarity check:

• N-H bonds are polar (ΔEN = |3.0 - 2.1| = 0.9)
• Asymmetric trigonal pyramidal geometry (lone pair)
• N-H dipoles do NOT cancel
• Polar molecule

IMFs present:

1. London Dispersion Forces (LDF) ✓

   • All molecules have LDF

2. Dipole-Dipole Forces ✓

   • NH₃ is polar
   • Permanent dipole moment

3. Hydrogen Bonding ✓

   • H bonded directly to N (nitrogen) ✓
   • Other NH₃ molecules have lone pair on N to accept H-bond ✓
   • Both requirements met!

Hydrogen bonding in NH₃:

• Each NH₃ can form up to 4 H-bonds
• 3 H atoms as donors
• 1 lone pair as acceptor

$H-N-H ... N-H$

Answer for (c): Hydrogen bonding, dipole-dipole forces, and London dispersion forces

Summary:

Expected boiling point order: NH₃ > HCl > CH₄

Actual values:

• CH₄: -161°C (weakest IMF)
• HCl: -85°C (dipole-dipole)
• NH₃: -33°C (H-bonding, strongest IMF)

Key takeaway:

• Nonpolar → LDF only
• Polar (no N-H/O-H/F-H) → DD + LDF
• Has N-H, O-H, or F-H → H-bonding + DD + LDF