Dipole-Dipole Interaction
Dipole-dipole interactions are intermolecular forces that occur between polar molecules, where the partially positive end (δ+) of one molecule is attracted to the partially negative end (δ–) of a neighbouring molecule. These forces are stronger than London dispersion forces but weaker than hydrogen bonds, typically ranging from 1–20 kJ/mol. They are responsible for the elevated boiling points of polar molecules such as HCl, SO₂, and acetone compared to nonpolar molecules of similar molecular weight.
Key Formula
U_dd = -(2μ₁²μ₂²) / (3 k_B T r⁶)
LaTeX: U_{\text{dd}} = -\frac{2\mu_1^2 \mu_2^2}{3k_B T r^6}
| Symbol | Meaning | Unit |
|---|---|---|
| μ₁, μ₂ | Dipole moments of the two molecules | C·m (or Debye) |
| k_B | Boltzmann constant (1.38 × 10⁻²³) | J/K |
| T | Temperature | K |
| r | Intermolecular distance | m |
| U_dd | Dipole-dipole interaction energy | J |
Worked Example
Problem
HCl (MW = 36.5 g/mol, dipole moment = 1.08 D) has a boiling point of –85 °C. Compare this to Ar (MW = 40 g/mol, nonpolar), which boils at –186 °C. Explain the difference.
Solution
Step 1: Ar is nonpolar; only London dispersion forces act between Ar atoms. Step 2: HCl is polar (1.08 D dipole moment); it experiences both London dispersion forces AND dipole-dipole interactions. Step 3: Despite HCl having slightly lower molecular weight than Ar (36.5 vs 40 g/mol), its boiling point is 101 °C higher. Step 4: The additional dipole-dipole attractions in HCl require more energy to overcome during boiling. Step 5: This demonstrates that polarity adds significant intermolecular attraction beyond what dispersion forces alone provide.
Answer
HCl boils at –85 °C vs Ar at –186 °C, a difference of 101 °C, primarily because HCl has additional dipole-dipole interactions (1.08 D) that Ar lacks, requiring more thermal energy to vaporise.
Dipole Moments and Boiling Points of Selected Polar Molecules
| Molecule | Dipole Moment (D) | MW (g/mol) | Boiling Point (°C) |
|---|---|---|---|
| HF | 1.91 | 20 | 19.5 |
| H₂O | 1.85 | 18 | 100 |
| HCl | 1.08 | 36.5 | –85 |
| SO₂ | 1.63 | 64 | –10 |
| Acetone (CH₃COCH₃) | 2.91 | 58 | 56 |
| Chloroform (CHCl₃) | 1.04 | 119.4 | 61 |
Interactive Tools
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Related Terms
Hydrogen Bond
A hydrogen bond is an attractive intermolecular or intramolecular force between a hydrogen atom covalently bonded to a highly electronegative atom (N, O, or F) and a lone pair on another electronegative atom. With strengths typically ranging from 5 to 40 kJ/mol, hydrogen bonds are much weaker than covalent bonds but far stronger than other van der Waals forces. They are responsible for water's unusually high boiling point, the structure of DNA base pairs, and the secondary and tertiary structures of proteins.
London Dispersion Forces
London dispersion forces (LDFs) are the weakest form of van der Waals forces, arising from temporary, instantaneous dipoles caused by the random fluctuation of electron density in any atom or molecule. Although individually very weak (0.1–40 kJ/mol), they are the only intermolecular force present in nonpolar molecules and noble gases, and become significant in large, polarisable molecules. LDFs increase with molecular size, surface area, and number of electrons, explaining why larger alkanes have higher boiling points than smaller ones.
Van der Waals Forces
Van der Waals forces are a collective term for weak, short-range intermolecular attractions that arise from transient or permanent electric dipoles, including London dispersion forces, dipole-dipole interactions, and dipole-induced dipole interactions. Named after Dutch physicist Johannes Diderik van der Waals, these forces explain deviations of real gases from ideal behaviour and govern properties such as boiling points, surface tension, and the adhesion of geckos to surfaces. Their strength scales with molecular size and polarisability, making them significant for large molecules and noble gases.
From Latin "dipole" — "di" (two) + "polus" (pole) — referring to molecules with two opposite charge centres. The term "interaction" describes the electrostatic attraction between these opposite poles. The concept was developed in the context of polar molecule behaviour in the early 20th century.