What Are Intermolecular Forces?
Intermolecular forces (IMFs) are the attractions and repulsions between neighboring molecules or atoms. Unlike intramolecular forces (the covalent and ionic bonds within molecules), IMFs are weaker attractions between separate molecules. Despite their relative weakness, IMFs govern the physical properties of all substances: boiling points, melting points, vapor pressure, viscosity, surface tension, and solubility.
There are three main types of intermolecular forces, in increasing strength: 1. London dispersion forces (also called Van der Waals forces or induced dipole–induced dipole forces) 2. Dipole–dipole forces 3. Hydrogen bonds
1. London Dispersion Forces
London dispersion forces (named after physicist Fritz London) are the weakest type of IMF, yet they are present in all molecules — polar or nonpolar, atomic or molecular. They are the only IMF in nonpolar substances like N₂, CH₄, and noble gases.
How They Arise
Electrons are in constant motion. At any given instant, a nonpolar molecule like Cl₂ may have a slightly uneven electron distribution — an instantaneous dipole. This fleeting dipole induces a matching dipole in a neighboring molecule (an induced dipole). The two instantaneous dipoles attract each other momentarily.
Factors Affecting London Dispersion Strength
- Number of electrons (polarizability): More electrons = larger, more easily distorted electron cloud = stronger instantaneous dipoles = stronger London forces.
- Molecular size and shape: Larger, elongated molecules (like n-pentane) have stronger London forces than compact, spherical molecules (like neopentane) with the same formula.
Evidence from boiling points (nonpolar noble gases):
| Noble Gas | Electrons | Boiling Point |
|---|---|---|
| He | 2 | −269°C |
| Ne | 10 | −246°C |
| Ar | 18 | −186°C |
| Kr | 36 | −153°C |
| Xe | 54 | −108°C |
As the number of electrons increases, London forces strengthen and more energy is required to vaporize the substance.
2. Dipole–Dipole Forces
Dipole–dipole forces occur between polar molecules. A polar molecule has a permanent dipole moment — one end is δ+ and the other is δ−. Neighboring polar molecules align so that opposite partial charges attract each other.
These forces are stronger than London forces for molecules of similar size. They operate over short distances and are highly directional.
Example: HCl vs. F₂
Both HCl and F₂ have similar molecular masses (~36–38 g/mol), but: - HCl (polar, has dipole–dipole forces): boiling point = −85°C - F₂ (nonpolar, only London forces): boiling point = −188°C
The dipole–dipole interactions in HCl require more energy to overcome → higher boiling point.
Keesom Interactions
The thermal-average alignment of dipoles (called Keesom interactions) contributes to dipole–dipole forces. At higher temperatures, thermal motion disrupts the alignment, so dipole–dipole forces weaken with increasing temperature more than London forces do.
3. Hydrogen Bonds
Hydrogen bonds are a special, unusually strong type of dipole–dipole interaction that occurs when hydrogen is bonded directly to fluorine (F), oxygen (O), or nitrogen (N) — the three most electronegative elements.
When H is attached to F, O, or N: - The H becomes strongly δ+ (very electron-poor) - The F, O, or N carries a large δ− - The small size of H allows it to approach lone pairs on neighboring F, O, or N atoms very closely
The result is an unusually strong electrostatic attraction between the δ+ H and a lone pair on an adjacent molecule: the hydrogen bond (represented as A–H···B, where ··· is the hydrogen bond).
Why F, O, N Only?
Three conditions must be met for a true hydrogen bond: 1. H must be bonded to F, O, or N 2. The neighboring atom must be F, O, or N (the acceptor) 3. The geometry must allow H to approach the lone pair
C–H bonds are too weakly polar to form significant hydrogen bonds. Heavier elements (Cl, S) are too large for H to approach their lone pairs closely.
Evidence: The Anomalous Properties of Water
Water (H₂O, MW = 18 g/mol) boils at 100°C. Based on the trend in group 16 hydrides (H₂S boils at −60°C, H₂Se at −41°C, H₂Te at −2°C), water should boil around −80°C. Instead, it boils ~180°C higher — a direct consequence of its strong hydrogen bonding network.
Each water molecule can form up to 4 hydrogen bonds (2 as donor through its O–H bonds, 2 as acceptor through its lone pairs), creating a highly cohesive network.
Hydrogen Bonds in Biology
Hydrogen bonds are central to life: - DNA double helix: Held together by H-bonds between complementary base pairs (A–T: 2 H-bonds; G–C: 3 H-bonds) - Protein folding: α-helices and β-sheets are stabilized by H-bonds between backbone N–H and C=O groups - Enzyme active sites: Precise H-bonding geometries position substrates for catalysis - Cellulose: H-bonds between chains give plant cell walls their rigidity
Comparing IMF Strengths
| IMF Type | Present In | Relative Strength | Range |
|---|---|---|---|
| London dispersion | All molecules | Weakest | Very short |
| Dipole–dipole | Polar molecules | Medium | Short |
| Hydrogen bonds | H bonded to F, O, N | Strongest (of IMF) | Very short |
Note: All IMFs are much weaker than intramolecular covalent bonds. IMF energies are typically 1–40 kJ/mol, while covalent bonds are 150–1000 kJ/mol.
Physical Properties Governed by IMFs
Boiling Points
More energy is needed to vaporize a substance with stronger IMFs. Higher IMFs → higher boiling point.
Compare: - CH₄ (only London forces): bp = −161°C - CH₃Cl (dipole–dipole + London): bp = −24°C - CH₃OH (H-bonding + London): bp = 65°C
Viscosity
Stronger IMFs → greater resistance to flow → higher viscosity. Honey is more viscous than water (despite both being hydrogen-bonding liquids) largely because of its sugar molecules' extensive H-bonding network and molecular entanglement.
Surface Tension
Stronger IMFs → molecules at the surface are pulled more strongly inward → higher surface tension. Water's high surface tension (72 mN/m) allows insects to walk on water and water droplets to bead up.
Solubility
"Like dissolves like" — substances with similar IMFs dissolve in each other. Polar solvents dissolve polar solutes via dipole–dipole interactions; nonpolar solvents dissolve nonpolar solutes via London forces.
Ion–Dipole Forces (Bonus)
In solutions of ionic compounds, there is also a fourth important force: ion–dipole interactions. When NaCl dissolves in water, Na⁺ ions attract the δ− oxygen of water molecules, and Cl⁻ ions attract the δ+ hydrogen atoms. These ion–dipole forces are generally the strongest of all non-covalent interactions and are the driving force behind ionic compound dissolution.