Master AP Chemistry Topic 3.1 with this free practice quiz. Test your skills on london dispersion forces, polarizability, dipole dipole forces, and hydrogen bonding to prepare for your AP exam.
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Introduction to Intermolecular and Interparticle Forces
Intermolecular forces are attractions between separate particles, such as molecules. Interparticle forces is a broader term that also includes attractions between ions and molecules. These forces influence the physical properties of substances, including boiling point, melting point, viscosity, surface tension, vapor pressure, and solubility.
Intermolecular forces are different from intramolecular bonds, which hold atoms together within a particle. For example, the O–H bonds within a water molecule are covalent bonds, while the attractions between neighboring water molecules are intermolecular forces.
During boiling, water molecules separate from one another, overcoming intermolecular attractions without breaking their O–H covalent bonds. This distinction explains why physical changes usually involve overcoming attractions between particles rather than breaking the chemical bonds within them.
The strength of intermolecular forces depends on factors such as molecular polarity, electron distribution, particle size, and charge. AP Chemistry focuses on five important areas: London dispersion forces, dipole-dipole interactions, hydrogen bonding, ion-dipole forces, and their effects on macroscopic properties.
1. London Dispersion Forces
London dispersion forces arise from temporary changes in electron distribution. Electrons are constantly moving, so the electron cloud around an atom or molecule can become uneven for a moment. This creates a temporary dipole, which can induce a dipole in a neighboring particle. The resulting attraction is called a London dispersion force.
London dispersion forces occur between all atoms and molecules, including nonpolar substances. They are the only type of intermolecular force present in a pure sample of a nonpolar molecular substance, such as CH₄ or I₂.
Factors affecting dispersion forces
The strength of London dispersion forces depends mainly on polarizability, which describes how easily an electron cloud can be distorted.
Larger atoms and molecules generally have more electrons and more easily distorted electron clouds. Their dispersion forces are therefore usually stronger.
For example, the noble gases show increasing boiling points down Group 18:
He < Ne < Ar < Kr < Xe
As the atoms become larger and contain more electrons, their dispersion forces increase. More energy is needed to separate the atoms, so their boiling points rise.
Molecular shape also matters. Long, relatively unbranched molecules can have greater contact with neighboring molecules than compact, highly branched molecules of similar molar mass. Greater contact can strengthen dispersion attractions and increase boiling point.
For example, straight-chain pentane generally has a higher boiling point than its more compact, branched isomer neopentane, despite both having the molecular formula C₅H₁₂.
2. Dipole-Dipole Interactions
Dipole-dipole interactions occur between polar molecules. A polar molecule has an uneven distribution of electron density, giving it a partially positive end (δ+) and a partially negative end (δ−).
When polar molecules approach one another, the partially positive end of one molecule is attracted to the partially negative end of another. These attractions influence how much energy is required to separate the molecules.
Hydrogen chloride, HCl, is an example of a polar molecule. The chlorine atom attracts bonding electrons more strongly than hydrogen, so HCl has a permanent dipole. Neighboring HCl molecules experience dipole-dipole attractions.
The strength of these interactions depends on the magnitude of the molecular dipoles and the relative orientations of the molecules. Favorable orientations place oppositely charged ends near one another.
All polar molecules also experience London dispersion forces. Therefore, a pure sample of HCl has both dispersion forces and dipole-dipole interactions.
When comparing similar substances, stronger dipole-dipole attractions can lead to higher boiling points because more energy is required to separate the molecules. However, molar mass, molecular shape, and dispersion forces must also be considered before predicting which substance has the higher boiling point.
3. Hydrogen Bonding
Hydrogen bonding is a particularly strong type of dipole-dipole interaction. It occurs when hydrogen is covalently bonded to nitrogen, oxygen, or fluorine and is attracted to a lone pair of electrons on a nitrogen, oxygen, or fluorine atom in a neighboring particle.
The three important bonding arrangements are N–H, O–H, and F–H.
Hydrogen bonding is especially strong because nitrogen, oxygen, and fluorine are highly electronegative. They draw electron density away from hydrogen, leaving it strongly partially positive. This hydrogen can then be attracted to a lone pair on a nearby molecule.
Water provides a familiar example. The hydrogen atom of one H₂O molecule can form a hydrogen bond with the oxygen atom of another H₂O molecule. Each water molecule can participate in several hydrogen bonds, producing an extensive network of attractions.
Hydrogen bonding helps explain why water has an unusually high boiling point for its small molar mass. Consider the boiling points of the Group 16 hydrides:
H₂S boils at approximately −60 °C, while H₂O boils at 100 °C at standard atmospheric pressure.
Water molecules form strong hydrogen bonds, whereas H₂S does not form comparable hydrogen bonding. Substantially more energy is required to separate water molecules into the gas phase.
Hydrogen bonding also occurs between molecules such as NH₃ and HF. It can occur within a single molecule when its structure brings a suitable hydrogen and an appropriate lone pair close together, although AP Chemistry examples commonly emphasize attractions between separate molecules.
Hydrogen bonding influences boiling points, viscosity, surface tension, and solubility. It is also important in biological molecules, including the interactions that help stabilize proteins and DNA.
4. Ion-Dipole Forces
Ion-dipole forces are attractions between an ion and a polar molecule. They are particularly important when ionic compounds dissolve in polar solvents such as water.
Water has a permanent dipole. Its oxygen end is partially negative, while its hydrogen ends are partially positive. Consequently, water molecules orient themselves differently around positive and negative ions.
When sodium chloride dissolves in water, the ionic lattice separates into hydrated ions:
NaCl(s) → Na⁺(aq) + Cl⁻(aq)
Around Na⁺, the partially negative oxygen ends of water molecules point toward the cation. Around Cl⁻, the partially positive hydrogen ends point toward the anion. These attractions between ions and water molecules are ion-dipole forces.
The process is called hydration when water is the solvent. The notation (aq) indicates that an ion is dissolved in water and surrounded by solvent molecules.
Ion-dipole attractions help stabilize separated ions in solution. Whether an ionic compound dissolves readily depends on the balance between the energy required to separate its ions in the solid lattice and the energy released when the ions interact with the solvent. Entropy changes also contribute to the overall favorability of dissolution.
Ion-dipole forces generally become stronger as the ionic charge increases and the distance between the ion and the relevant dipole decreases. A small ion with a high charge density can interact strongly with polar molecules.
These interactions explain why water dissolves many ionic substances, although not every ionic compound is highly soluble. The lattice energy and other thermodynamic factors still matter.
5. Comparing Intermolecular and Interparticle Forces
The four forces differ in the particles involved and the origins of their attractions.
London dispersion forces occur between all atoms and molecules and result from temporary and induced dipoles.
Dipole-dipole interactions occur between permanent dipoles in polar molecules.
Hydrogen bonding is a strong, specific type of dipole-dipole interaction involving hydrogen bonded to N, O, or F and a suitable lone pair on a neighboring particle.
Ion-dipole forces occur between ions and polar molecules, especially during the dissolution of ionic compounds in polar solvents.
For particles of comparable size and structure, these attractions are often considered in the approximate order:
London dispersion < dipole-dipole < hydrogen bonding < ion-dipole
This is a useful guide rather than a universal ranking. Actual interaction strengths depend on particle size, charge, polarity, distance, and orientation. Large nonpolar molecules can have very strong dispersion forces, and ion-dipole interactions vary considerably with ionic charge and solvent structure.
When analyzing a substance, identify every relevant force rather than assigning only one. For example, liquid water exhibits dispersion forces, dipole-dipole interactions, and hydrogen bonding.
6. How Interparticle Forces Affect Macroscopic Properties
Macroscopic properties are observable properties of a bulk sample. They reflect the combined behavior of enormous numbers of particles and the attractions between them.
Boiling point and vapor pressure
Boiling occurs when a liquid’s vapor pressure equals the external pressure. Molecules must escape from the liquid into the gas phase, separating from neighboring particles.
At the same temperature, a liquid with stronger intermolecular attractions generally has a lower vapor pressure because fewer molecules can escape into the gas phase. More energy is usually needed to boil it, resulting in a higher boiling point when comparable substances are considered at the same external pressure.
For example, ethanol has a higher boiling point than dimethyl ether, even though both have the molecular formula C₂H₆O. Ethanol contains an O–H bond and can form hydrogen bonds with neighboring ethanol molecules. Dimethyl ether cannot donate hydrogen bonds in a pure sample, so its intermolecular attractions differ.
Viscosity
Viscosity measures a fluid’s resistance to flow. Stronger intermolecular attractions can make molecules more difficult to move past one another, increasing viscosity. Molecular shape and the extent of molecular entanglement can also influence viscosity.
For example, glycerol has three O–H groups per molecule and can form extensive hydrogen-bonding interactions. Its liquid is much more viscous than many small molecular liquids.
Surface tension
Surface tension is the energy required to increase the surface area of a liquid per unit area, or equivalently, a measure of the tendency of a liquid surface to contract.
Molecules at the surface have fewer neighboring liquid molecules than molecules in the bulk. Attractive forces therefore create a net energetic cost when additional surface is formed. Strong cohesive attractions generally produce higher surface tension.
Water’s hydrogen bonding contributes to its relatively high surface tension, allowing small insects to move across its surface under suitable conditions.
Melting point and phase changes
Melting requires particles in a solid to move from an ordered structure into a liquid. Strong attractions often increase melting points, but melting-point trends are less straightforward than boiling-point trends because crystal structure, symmetry, and packing also matter.
During melting and boiling, energy is used to overcome or reorganize interparticle attractions. During these phase changes, the temperature remains constant while the phase transition occurs at constant pressure because the energy changes the arrangement and separation of particles rather than their average kinetic energy.
7. Applying Interparticle Forces to Solubility
Solubility describes how much solute dissolves in a solvent under specified conditions. Dissolution depends on the interactions among solute particles, solvent particles, and newly formed solute-solvent interactions.
A useful principle is that substances tend to dissolve when the new interactions and entropy changes make the process thermodynamically favorable.
Polar substances often dissolve well in polar solvents because they can form favorable dipole-dipole, hydrogen-bonding, or ion-dipole interactions. Nonpolar substances often dissolve better in nonpolar solvents because dispersion forces can provide favorable interactions without a major mismatch in polarity.
For example, ethanol mixes readily with water because its O–H group can form hydrogen bonds with water. Hexane, a nonpolar hydrocarbon, does not mix appreciably with water because mixing does not provide sufficiently favorable interactions to compensate for the changes involved in separating and reorganizing the particles.
In AP Chemistry, solubility predictions should consider molecular structure and the types of attractions possible between the solute and solvent, rather than relying on polarity alone. Ionic compounds also require consideration of lattice energy and hydration.
These principles connect microscopic particle attractions to observable properties, allowing molecular structure and interparticle forces to explain differences in boiling point, vapor pressure, viscosity, surface tension, and solubility.