10.1 Types of intermolecular forces

Intermolecular Forces: Types and Strengths

The Main Types of Intermolecular Forces

Three categories of intermolecular forces govern how molecules attract one another: London dispersion forces, dipole-dipole interactions, and hydrogen bonding. Each arises from a different physical origin, but all are electrostatic in nature.

• London dispersion forces (also called induced dipole–induced dipole interactions) are present between all molecules, including nonpolar ones. They arise from temporary fluctuations in electron density that create instantaneous dipoles; these fleeting dipoles then induce dipoles in neighboring molecules, producing a net attraction (e.g., Ar, $\text{CH}_4$, $\text{N}_2$).
• Dipole-dipole interactions occur between polar molecules. The partially positive end ($\delta^+$) of one molecule is attracted to the partially negative end ($\delta^-$) of another (e.g., $\text{HCl}$, $\text{SO}_2$).
• Hydrogen bonding is a particularly strong subset of dipole-dipole interaction. It occurs when a hydrogen atom bonded to a highly electronegative atom (N, O, or F) interacts with a lone pair on another N, O, or F atom. The small radius of hydrogen allows an unusually close approach, making the interaction stronger and more directional than a typical dipole-dipole force (e.g., $\text{H}_2\text{O}$, $\text{NH}_3$, $\text{HF}$).

Relative Strengths of Intermolecular Forces

As a rough ranking from weakest to strongest:

$\text{London dispersion} < \text{dipole-dipole} < \text{hydrogen bonding}$

This ranking holds for molecules of comparable size, but there are important nuances:

• London dispersion forces scale with polarizability. Larger molecules with more electrons have a more easily distorted electron cloud, so their dispersion forces can be substantial. For instance, $\text{I}_2$ (a nonpolar molecule) is a solid at room temperature because its large electron cloud produces strong dispersion forces, while $\text{Cl}_2$ is a gas.
• Dipole-dipole strength depends on dipole moment magnitude. Molecules with larger permanent dipole moments experience stronger attractions. $\text{HF}$ ($\mu = 1.83 \text{ D}$) has a larger dipole moment than $\text{HCl}$ ($\mu = 1.08 \text{ D}$), contributing to stronger dipole-dipole interactions.
• Hydrogen bonds are typically the strongest of the three because the high electronegativity of N, O, or F concentrates charge on the hydrogen, and hydrogen's lack of inner-shell electrons leaves it partially unshielded. This combination produces interactions roughly 5–30 kJ/mol in strength, compared to ~0.05–40 kJ/mol for dispersion forces (range depends heavily on molecular size) and ~5–25 kJ/mol for ordinary dipole-dipole forces.

Origin and Nature of Intermolecular Forces

Dipole-Dipole Interactions

Polar molecules have a permanent dipole moment because bonded atoms differ in electronegativity, pulling electron density unevenly across the bond. In $\text{HCl}$, chlorine's higher electronegativity draws electron density toward itself, giving the molecule a permanent $\delta^+$ (H end) and $\delta^-$ (Cl end).

When two polar molecules approach each other, they orient so that opposite partial charges face one another. The interaction energy between two dipoles falls off as $\frac{1}{r^3}$ for fixed orientations, or as $\frac{1}{r^6}$ when thermal averaging is included (the Keesom interaction). This steep distance dependence means dipole-dipole forces are most significant when molecules are close together in the liquid or solid phase.

Comparing $\text{HCl}$ ($\mu = 1.08 \text{ D}$, bp −85 °C) with $\text{HBr}$ ($\mu = 0.82 \text{ D}$, bp −67 °C) illustrates a subtlety: $\text{HBr}$ actually boils higher despite its smaller dipole moment, because its larger electron cloud produces stronger London dispersion forces that more than compensate. This is a good reminder that you can't look at one type of force in isolation.

London Dispersion Forces

At any instant, the electrons in a molecule may be distributed asymmetrically, creating a transient dipole. This instantaneous dipole induces a complementary dipole in a neighboring molecule, and the two attract each other. Averaged over time, the net effect is always attractive.

Key factors controlling dispersion force strength:

• Number of electrons / molar mass. More electrons means a larger, more easily distorted electron cloud. $\text{I}_2$ (254 g/mol, solid at 25 °C) vs. $\text{F}_2$ (38 g/mol, gas at 25 °C) is a dramatic example.
• Molecular shape. Elongated molecules have more surface area for contact than compact, spherical ones. n-Pentane (bp 36 °C) boils higher than neopentane (bp 9.5 °C) even though they share the same molecular formula ($\text{C}_5\text{H}_{12}$).

Dispersion forces are the only intermolecular force available to nonpolar species like noble gases (He, Ne, Ar) and homonuclear diatomics ($\text{N}_2$, $\text{O}_2$). They are also present in polar molecules, layered on top of dipole-dipole and hydrogen-bonding contributions.

Hydrogen Bonding

Hydrogen bonding requires a specific arrangement: a hydrogen atom covalently bonded to N, O, or F (the donor) interacting with a lone pair on another N, O, or F atom (the acceptor).

Why is this interaction so much stronger than an ordinary dipole-dipole force?

1. N, O, and F are among the most electronegative elements, so the $\text{X–H}$ bond is highly polar.
2. Hydrogen has no inner-shell electrons. Once its bonding electrons are pulled toward the electronegative atom, the remaining positive charge is barely shielded, allowing an unusually close approach to the acceptor's lone pair.
3. The interaction is strongly directional, typically forming along the axis of the lone pair, which gives it partial covalent character.

Water is the classic example. Each $\text{H}_2\text{O}$ molecule can donate two hydrogen bonds (one per H) and accept two (one per lone pair), creating an extensive three-dimensional network. This network explains water's anomalously high boiling point (100 °C) compared to $\text{H}_2\text{S}$ (−60 °C), which cannot hydrogen bond despite being heavier.

Hydrogen bonding also underpins DNA base pairing: guanine–cytosine pairs form three hydrogen bonds, while adenine–thymine pairs form two, making G–C pairs more thermally stable.

Intermolecular Forces: Impact on Properties

Physical Properties

Intermolecular forces directly determine melting point, boiling point, vapor pressure, viscosity, and surface tension. The logic is straightforward: stronger attractions between molecules mean you need more energy to pull them apart.

• Boiling and melting points increase with stronger intermolecular forces. Water ($\text{H}_2\text{O}$, bp 100 °C) boils far higher than methane ($\text{CH}_4$, bp −161 °C) despite its lower molar mass, because water has extensive hydrogen bonding while methane relies solely on weak dispersion forces.
• Vapor pressure decreases as intermolecular forces strengthen, because fewer molecules have enough kinetic energy to escape into the gas phase.
• Viscosity and surface tension both increase with stronger intermolecular forces. Glycerol ($\text{C}_3\text{H}_8\text{O}_3$), which can form up to three hydrogen bonds per molecule, is far more viscous than ethanol ($\text{C}_2\text{H}_5\text{OH}$), which can form only one donor hydrogen bond.

Solubility and Miscibility

The guiding principle is "like dissolves like": substances with similar intermolecular forces tend to be mutually soluble.

• Polar solutes in polar solvents. $\text{NaCl}$ dissolves readily in water because strong ion-dipole interactions between $\text{Na}^+$/$\text{Cl}^-$ and $\text{H}_2\text{O}$ replace the lattice energy of the crystal. In nonpolar hexane ($\text{C}_6\text{H}_{14}$), no comparable interactions exist, so $\text{NaCl}$ is essentially insoluble.
• Nonpolar solutes in nonpolar solvents. Hexane and benzene are miscible because both interact primarily through dispersion forces of similar magnitude.
• Miscibility of liquids follows the same logic. Ethanol and water are fully miscible because ethanol's $\text{OH}$ group hydrogen bonds with water. Oil (long nonpolar hydrocarbon chains) and water are immiscible because the strong hydrogen-bond network of water would have to be disrupted without adequate replacement interactions.

Biological Systems

Intermolecular forces are central to the structure and function of biological macromolecules.

• Protein secondary structure. The $\alpha$-helix is stabilized by hydrogen bonds between the $\text{C=O}$ of one amino acid and the $\text{N–H}$ four residues later along the chain. $\beta$-sheets are held together by hydrogen bonds between adjacent strands.
• DNA double helix. Complementary base pairing relies on hydrogen bonds (G–C: 3 bonds, A–T: 2 bonds), providing both stability and specificity for genetic information storage.
• The hydrophobic effect. When nonpolar molecules are introduced into water, they disrupt the hydrogen-bond network. Water molecules reorganize around the nonpolar surface, losing entropy. To minimize this disruption, nonpolar groups aggregate together. This drives protein folding (burying hydrophobic residues in the interior), micelle formation by surfactants, and the assembly of lipid bilayers in cell membranes.

Intermolecular Forces: Prediction and Interpretation

Predicting Relative Strengths

To predict which substance has stronger intermolecular forces, work through these steps:

1. Identify the types of forces present. Can the molecule hydrogen bond (H bonded to N, O, or F)? Is it polar (dipole-dipole)? All molecules have London dispersion forces.
2. Compare hydrogen-bonding capability. If one molecule can hydrogen bond and the other cannot, the hydrogen-bonding molecule usually has stronger overall intermolecular forces (assuming similar size).
3. Compare polarity. Among non-hydrogen-bonding molecules, a larger dipole moment generally means stronger dipole-dipole interactions.
4. Compare size and polarizability. For molecules with similar polarity, the larger molecule will have stronger dispersion forces.

Example: Ethanol ($\text{CH}_3\text{CH}_2\text{OH}$) vs. ethane ($\text{CH}_3\text{CH}_3$). Ethanol can form hydrogen bonds; ethane cannot. Despite similar molar masses (~46 vs. 30 g/mol), ethanol boils at 78 °C while ethane boils at −89 °C.

Explaining Differences in Physical Properties

Once you've identified the dominant intermolecular forces, you can explain property differences quantitatively.

• Water ($\text{H}_2\text{O}$, 18 g/mol, bp 100 °C) vs. methane ($\text{CH}_4$, 16 g/mol, bp −161 °C): nearly the same molar mass, but a 261 °C difference in boiling point. Water's extensive hydrogen-bond network requires far more energy to disrupt.
• Ethanol vs. glycerol: glycerol has three $\text{OH}$ groups and can form a denser hydrogen-bond network, resulting in much higher viscosity (about 1,500 mPa·s vs. 1.1 mPa·s at 20 °C).

Predicting Solubility and Miscibility

Apply the "like dissolves like" principle systematically:

1. Determine the dominant intermolecular forces in the solute.
2. Determine the dominant intermolecular forces in the solvent.
3. If they match (both polar, both nonpolar, or both capable of hydrogen bonding), predict good solubility. If they differ significantly, predict poor solubility.

Example: Will caffeine dissolve better in water or in dichloromethane ($\text{CH}_2\text{Cl}_2$)? Caffeine is a moderately polar molecule with hydrogen-bond acceptor sites. Dichloromethane is polar enough to interact with caffeine's polar groups, and caffeine's aromatic rings also interact favorably with the organic solvent. In practice, caffeine is more soluble in dichloromethane than in water, which is why $\text{CH}_2\text{Cl}_2$ is used to extract caffeine from aqueous coffee solutions.

Interpreting Behavior of Mixtures and Solutions

Real-world phenomena often involve a competition between different intermolecular forces.

• Micelle formation. Surfactant molecules have a polar head and a nonpolar tail. In water, the tails cluster inward to avoid disrupting water's hydrogen-bond network, while the polar heads face outward and interact with water. The result is a spherical micelle, driven by the hydrophobic effect.
• Liquid-liquid extraction. Separating a polar compound from a nonpolar one works because each preferentially dissolves in the solvent whose intermolecular forces match its own. Shaking a mixture of caffeine (polar) and oils (nonpolar) with water and dichloromethane partitions the components into different layers based on their intermolecular interactions.