Premed · Premed · General Chemistry 1
Lecture 16: Intermolecular Forces
General Chemistry I
Learning Objectives
By the end of this lecture, students will be able to:
- Distinguish between intramolecular forces (chemical bonds) and intermolecular forces (IMFs)
- Identify and rank the types of intermolecular forces: London dispersion, dipole-dipole, and hydrogen bonding
- Explain the origin of each type of IMF at the molecular level
- Predict the types and relative strengths of IMFs present in a given substance
- Relate intermolecular forces to physical properties such as boiling point, viscosity, and surface tension
Lecture Content
I. Intramolecular vs. Intermolecular Forces
Intramolecular forces are the forces within a molecule -- covalent and ionic bonds that hold atoms together. These are strong, typically ranging from 100 to 1000 kJ/mol. Intermolecular forces (IMFs) are the forces between molecules that hold them together in condensed phases (liquids and solids). These are much weaker, typically ranging from 0.05 to 40 kJ/mol. IMFs determine physical properties such as boiling point, melting point, viscosity, surface tension, and solubility, while intramolecular forces govern chemical properties. A critical distinction to keep in mind: when a molecular substance boils, it is the intermolecular forces that are overcome, not the covalent bonds within the molecules. Boiling water, for instance, breaks hydrogen bonds between H2O molecules while leaving the O-H covalent bonds completely intact.
II. Types of Intermolecular Forces (Weakest to Strongest)
A. London Dispersion Forces (LDF) -- Present in ALL Molecules
London dispersion forces, also called van der Waals forces or induced dipole-induced dipole forces, arise from temporary, instantaneous dipoles that form due to random fluctuations in a molecule's electron distribution. At any given instant, the electron cloud may be asymmetric, creating a momentary dipole that then induces a dipole in a neighboring molecule. These induced dipoles attract each other. London dispersion forces are present in every substance, whether polar or nonpolar.
The strength of LDF depends primarily on polarizability -- the ease with which an electron cloud can be distorted. Molecules with more electrons (higher molar mass) have larger, more polarizable electron clouds and therefore stronger London forces. Molecular shape also matters: elongated molecules have greater surface area for contact, producing stronger LDF. This is why n-pentane (a straight-chain molecule) has a higher boiling point than neopentane (a compact, spherical molecule) despite having the same molecular formula. For nonpolar molecules such as He, N2, CH4, CCl4, and I2, London dispersion forces are the only intermolecular force present. For large nonpolar molecules, these forces can become quite substantial -- iodine (I2), for instance, is a solid at room temperature entirely due to strong London forces.
B. Dipole-Dipole Forces
Dipole-dipole forces arise from the attraction between permanent dipoles in polar molecules. The positive end (delta+) of one molecule attracts the negative end (delta-) of a neighboring molecule. These forces are present only in polar molecules and their strength depends on the magnitude of the molecular dipole moment. For molecules of similar size and molar mass, dipole-dipole forces are stronger than London forces alone. Comparing CH3Cl (polar, boiling point -24 C) with C2H6 (nonpolar, similar molecular weight, boiling point -89 C) illustrates this difference nicely. In the condensed phase, polar molecules tend to orient themselves so that opposite partial charges are adjacent.
C. Hydrogen Bonding
Hydrogen bonding is a special, unusually strong form of dipole-dipole interaction that occurs when hydrogen is bonded directly to fluorine, oxygen, or nitrogen -- small, highly electronegative atoms that possess lone pairs. The partially positive hydrogen (delta+) on one molecule is attracted to a lone pair on the F, O, or N of another molecule. This interaction is written as D-H...A, where D is the donor atom bonded to hydrogen and A is the acceptor atom with the lone pair; both must be F, O, or N. Hydrogen bonds are considerably stronger than typical dipole-dipole forces, approximately 10 to 40 kJ/mol.
The consequences of hydrogen bonding are far-reaching. Water, hydrogen fluoride, and ammonia all have anomalously high boiling points compared to heavier analogs within their respective groups. Water's high specific heat, high heat of vaporization, and existence as a liquid at room temperature all stem from extensive hydrogen bonding. In biology, hydrogen bonds are essential for DNA base pairing, protein folding, and enzyme-substrate interactions.
<image>A three-panel illustration of the three types of intermolecular forces. Panel A (London Dispersion Forces): two nonpolar molecules (e.g., Ar atoms) side by side. A sequence shows: (1) symmetrical electron clouds, (2) instantaneous dipole forms in one atom with delta+ and delta- regions, (3) the neighboring atom's electron cloud is distorted (induced dipole), and the two attract. Panel B (Dipole-Dipole): two HCl molecules aligned with the delta+ end of one near the delta- end of the other, with an arrow showing the attractive force. Multiple molecules shown in a condensed-phase arrangement with alternating orientations. Panel C (Hydrogen Bonding): three water molecules connected by hydrogen bonds shown as dashed lines from H (delta+) on one molecule to O (delta-) on another. Lone pairs on O are shown explicitly. Bond lengths are labeled: O-H covalent bond ~96 pm, O...H hydrogen bond ~197 pm.</image>
D. Ion-Dipole Forces
Although not strictly an intermolecular force between neutral molecules, ion-dipole forces play a crucial role in solution chemistry. They arise from the attraction between an ion and the partial charge of a polar molecule. Sodium ions in water, for example, are surrounded by water molecules with their oxygen atoms (delta-) oriented toward the cation. Ion-dipole forces are the strongest of these intermolecular-type interactions and are directly responsible for the dissolution of ionic compounds in water through the hydration of ions.
III. Ranking Intermolecular Forces
The general ranking of intermolecular force strength, from strongest to weakest, is: ion-dipole > hydrogen bonding > dipole-dipole > London dispersion. However, London dispersion forces can become very strong for large molecules and may exceed dipole-dipole forces in such cases. The dominant intermolecular force in a substance ultimately determines its physical properties. A systematic strategy for identifying IMFs in any substance begins with asking: Is it ionic (ionic bonding, even stronger than IMFs)? Is it polar (dipole-dipole forces are present)? Does it have H bonded to F, O, or N (hydrogen bonding is present)? And finally, all molecules possess London dispersion forces.
IV. Relating IMFs to Physical Properties
A. Boiling Point and Melting Point
Stronger intermolecular forces require more energy to overcome, leading to higher boiling and melting points. Among molecules of similar size, boiling point rises with increasing IMF strength. The progression N2 (LDF only, bp -196 C) < CO (weak dipole, bp -192 C) < HCl (dipole-dipole, bp -85 C) < H2O (hydrogen bonding, bp 100 C) illustrates this beautifully. Within a homologous series where the same type of IMF dominates, boiling point increases with molar mass as London forces grow stronger: CH4 < C2H6 < C3H8 < C4H10.
B. Surface Tension
Surface tension is the tendency of a liquid surface to minimize its area, driven by the inward pull on surface molecules from their neighbors below and beside them. Stronger IMFs produce higher surface tension. Water's high surface tension is a direct consequence of its extensive hydrogen bonding network.
C. Viscosity
Viscosity measures a liquid's resistance to flow. Stronger intermolecular forces and larger molecular size both increase viscosity. Viscosity decreases with increasing temperature as molecules gain kinetic energy and overcome attractive forces more readily.
D. Capillary Action
Capillary action is the movement of a liquid through narrow spaces against gravity, driven by the competition between adhesive forces (attraction between the liquid and the surface) and cohesive forces (attraction between liquid molecules). Water rises in glass tubes because its adhesion to the glass surface (via hydrogen bonding with SiO2) exceeds its internal cohesion. Mercury in glass shows the opposite behavior: mercury's strong metallic cohesion exceeds its adhesion to glass, causing the meniscus to curve downward.
<image>A diagram comparing boiling points of hydrides across Groups 4A, 5A, 6A, and 7A of the periodic table. The x-axis shows period number (2 through 5) and the y-axis shows boiling point (in degrees C, ranging from -200 to +100). Four lines connect the hydrides: Group 4A (CH4, SiH4, GeH4, SnH4) -- shows steadily increasing BP with period. Group 5A (NH3, PH3, AsH3, SbH3) -- NH3 is anomalously high. Group 6A (H2O, H2S, H2Se, H2Te) -- H2O is dramatically anomalously high. Group 7A (HF, HCl, HBr, HI) -- HF is anomalously high. The anomalously high boiling points of H2O, HF, and NH3 are highlighted with annotations: "Hydrogen bonding causes anomalously high boiling points for the Period 2 hydrides." The regular trend for Group 4A (no H-bonding) is contrasted as a baseline.</image>
V. Water: A Unique Substance
Water's extraordinary properties all trace back to extensive hydrogen bonding. Its boiling point, melting point, specific heat, and heat of vaporization are all anomalously high. Each water molecule can form up to four hydrogen bonds -- two through its donor H atoms and two through its acceptor lone pairs on oxygen. In ice, hydrogen bonds create an open hexagonal lattice that contains more empty space than liquid water, making ice less dense than the liquid. This unusual property is critical for aquatic life, as floating ice insulates the water below and prevents lakes from freezing solid. Water's high specific heat helps moderate climate and stabilize body temperature. And its high dielectric constant makes it an excellent solvent for ionic and polar substances.
VI. Predicting IMFs -- Practice Framework
To predict the intermolecular forces present in any molecule, begin by determining its geometry and polarity from its Lewis structure and VSEPR analysis. Then identify all applicable IMFs: all molecules have London dispersion forces (consider size and molar mass to assess their strength); polar molecules additionally exhibit dipole-dipole forces; and molecules with H-F, H-O, or H-N bonds also engage in hydrogen bonding. The dominant (strongest) IMF present determines the physical behavior. When comparing two molecules, first identify the strongest IMF in each. If both share the same dominant IMF, the one with greater molar mass will generally have the higher boiling point due to stronger London forces.
