AP Chemistry 3.1: Intermolecular Forces

• Intermolecular forces (IMFs) are attractive forces between separate particles, especially molecules. The prefix inter- means "between." This is different from intramolecular forces , which hold atoms together inside a molecule. For example, the O—H covalent bonds inside one H₂O molecule are intramolecular bonds, while attractions between two different H₂O molecules are intermolecular forces.

  • Intramolecular bonds are generally much stronger than intermolecular attractions. Boiling water does not normally break the O—H covalent bonds. It mainly separates H₂O molecules from one another by overcoming intermolecular attractions.

• Intermolecular forces strongly influence physical properties including boiling point, vapor pressure, evaporation rate, viscosity, surface tension, solubility, melting behavior, and physical state .

• The central idea for this entire topic is: the stronger the attractions between particles, the harder it is to separate those particles.

  • Stronger IMFs therefore generally produce higher boiling points , lower vapor pressures , slower evaporation rates , higher viscosities , and greater surface tensions .

  • The major attractions to know are London dispersion forces, dipole-dipole forces, hydrogen bonding, and ion-dipole forces . Ionic solids also involve strong ion-ion electrostatic attractions.

• London dispersion forces (LDFs) are present in every atom and molecule , whether the particle is polar or nonpolar. They arise because electrons are constantly moving. At any instant, the electron distribution may become temporarily uneven, producing an instantaneous dipole .

  • An instantaneous dipole can distort the electron cloud of a nearby particle. This creates an induced dipole in the neighboring particle. The temporary opposite partial charges attract each other.

  • Even noble gases such as Ne or Ar experience London dispersion forces. A nonpolar molecule such as CH₄ also experiences LDFs.

• Polarizability describes how easily a particle's electron cloud can be distorted. A more easily distorted electron cloud can create stronger temporary dipoles, producing stronger London dispersion forces.

  • Larger atoms and molecules generally contain more electrons and larger electron clouds, which are generally more polarizable.

  • Therefore, for similar types of particles, more electrons / larger electron cloud → stronger London dispersion forces .

  • For example, I₂ has far more electrons than F₂, so I₂ has much stronger dispersion forces. This helps explain why iodine can be a solid under conditions where fluorine is a gas.

• Molecular shape and surface area also affect dispersion forces. Long, relatively unbranched molecules can often touch neighboring molecules over a larger surface area than compact, highly branched molecules.

  • Greater surface contact means temporary dipoles can interact over a larger region, so greater molecular surface area can produce stronger LDFs .

  • This is why, among molecules with similar molar masses and functional groups, a less-branched molecule can sometimes have a higher boiling point than a more-branched isomer.

• Dipole-dipole forces occur between polar molecules . A polar molecule has a permanent uneven distribution of electron density.

  • The partially positive region, δ+, of one molecule is attracted to the partially negative region, δ−, of another.

  • To decide whether a molecule has dipole-dipole attractions, you must first determine whether the entire molecule is polar.

• Molecular polarity depends on both bond polarity and molecular geometry .

  • A molecule can contain polar bonds but still be nonpolar overall if the individual bond dipoles cancel.

  • CO₂ is a classic example. Each C=O bond is polar, but CO₂ is linear. The two equal bond dipoles point in opposite directions and cancel, so the whole molecule is nonpolar.

  • Because CO₂ is nonpolar, pure CO₂ does not experience permanent dipole-dipole attractions. It still experiences London dispersion forces.

• Hydrogen bonding is a particularly strong type of dipole-dipole attraction.

  • For standard AP Chemistry questions, look for hydrogen directly covalently bonded to N, O, or F .

  • The key bonds are therefore N—H, O—H, and F—H .

  • N, O, and F are especially electronegative and relatively small. When H is bonded to one of them, the H atom becomes strongly δ+, producing a strong attraction to a lone pair on N, O, or F of another particle.

  • H₂O can hydrogen-bond because it contains O—H bonds and oxygen has lone pairs.

  • NH₃ can hydrogen-bond because it contains N—H bonds and nitrogen has a lone pair.

  • HF can hydrogen-bond because it contains H—F.

  • CH₄ cannot hydrogen-bond with itself because its hydrogen atoms are bonded to carbon rather than N, O, or F.

• A common mistake is assuming that any molecule containing oxygen automatically hydrogen-bonds with itself. It does not.

  • For example, CH₃OCH₃ contains oxygen and is polar, but it has no O—H bond. Pure CH₃OCH₃ therefore has LDFs and dipole-dipole forces, but it cannot donate a hydrogen bond to another molecule of itself in the normal AP model.

  • However, a molecule with O or N lone pairs but no O—H or N—H bond can often still accept hydrogen bonds from another molecule that has an appropriate hydrogen-bond donor.

• Ion-dipole forces occur between an ion and a polar molecule.

  • These are especially important when ionic compounds dissolve in polar solvents such as water.

  • Water is polar: oxygen is δ− and hydrogen is δ+.

  • Around a positive ion such as Na⁺, the oxygen ends of nearby water molecules point toward the ion because Na⁺ attracts the δ− oxygen region.

  • Around a negative ion such as Cl⁻, the hydrogen ends point toward the ion because Cl⁻ attracts the δ+ hydrogen regions.

  • These ion-dipole attractions can help stabilize separated ions in solution.

• When asked for all intermolecular forces in a molecular substance, remember that LDFs are always present.

  • NH₃ therefore has London dispersion forces + dipole-dipole forces + hydrogen bonding .

  • HCl has London dispersion + dipole-dipole , but not normal hydrogen bonding.

  • CH₄ has London dispersion only .

• A rough conceptual ordering is often written as ion-ion > ion-dipole > hydrogen bonding > ordinary dipole-dipole > London dispersion , but you should not treat this as an absolute ranking for every substance .

• A very large, highly polarizable nonpolar molecule can have stronger total dispersion attractions than a very small polar molecule.

  • When AP Chemistry asks you to compare boiling points or vapor pressures, consider type of IMF, polarity, hydrogen-bonding capability, number of electrons, polarizability, molecular size, and molecular shape .

  • Khan Academy's current Unit 3 IMF material specifically covers London dispersion forces, dipole-dipole forces, hydrogen bonding, ion-dipole forces, vapor pressure, and using IMF reasoning to compare liquid properties.

• Evaporation occurs when molecules at the surface of a liquid have enough kinetic energy to overcome attractive forces and enter the gas phase.

  • Stronger IMFs make escape more difficult, so stronger IMFs → slower evaporation .

• Vapor pressure is the pressure produced by vapor above a liquid when the liquid and vapor reach dynamic equilibrium in a closed system.

  • At dynamic equilibrium, the rate of evaporation equals the rate of condensation . Individual molecules are still moving between phases, but the overall amounts remain approximately constant.

  • Strong IMFs hold particles in the liquid more strongly, so fewer particles escape into the gas phase.

  • Therefore stronger IMFs → lower vapor pressure .

  • Conversely, weak IMFs allow more particles to escape, producing higher vapor pressure .

  • Increasing temperature gives particles greater average kinetic energy. More particles can overcome their attractions, so temperature ↑ → vapor pressure ↑ .

• Boiling occurs when the liquid's vapor pressure equals the external pressure.

  • At normal atmospheric pressure, the temperature where this happens is called the normal boiling point .

  • Therefore, a liquid does not inherently "boil at 100°C." Water's normal boiling point is approximately 100°C at 1 atm, but changing the external pressure changes the boiling temperature.

  • At high altitude, atmospheric pressure is lower, so a liquid reaches the required vapor pressure at a lower temperature. Thus lower external pressure → lower boiling point .

  • Higher external pressure requires a higher vapor pressure before boiling begins, so higher external pressure → higher boiling point .

  • Because stronger IMFs require more energy to separate molecules, stronger IMF → higher boiling point .

• Viscosity is a liquid's resistance to flowing. Strong intermolecular attractions generally make it harder for molecules to slide past each other, giving greater viscosity.

  • Molecular size and shape can also influence viscosity. Long molecules can interact over larger areas and may become physically entangled.

• Surface tension comes from the net inward attractions experienced by molecules at a liquid's surface.

  • Molecules inside the liquid are attracted in many directions. Surface molecules have fewer neighboring particles above them, creating a net inward effect.

  • Stronger IMFs generally produce greater surface tension .

The most important liquid-property relationships to memorize are:

• Khan Academy's current liquid-property practice explicitly asks students to use intermolecular forces to compare boiling behavior.