Chemistry 2e · Liquids and Solids
Intermolecular Forces
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In 30 seconds
Inside a molecule, atoms are held together by intramolecular bonds — covalent bonds sharing electrons, or ionic bonds built from opposite charges. Between separate molecules, a different, weaker set of attractions operates: intermolecular forces (IMFs). These forces do not hold atoms together; they hold molecules near each other. They are the reason a gas condenses to a liquid, a liquid freezes to a solid, water beads on a waxed car, and oil and water refuse to mix.
This topic identifies the four main types of intermolecular forces — dispersion (London) forces, dipole A molecule (or bond) with separated partial positive and negative ends Full entry →–dipole attractions, hydrogen bonds, and ion–dipole forces — explains what determines their strength, and shows how they control physical properties like boiling point, vapor pressure, and solubility.
Why this matters
Nearly every "why" question about everyday matter is an IMF question: why water boils at 100 °C while methane boils at −162 °C, why a spider walks on a pond, why a paper towel soaks up a spill, why DNA's double helix stays zipped. In industry, IMFs decide which solvents dissolve which materials and how to separate mixtures by distillation. In biology, hydrogen bonds stabilize proteins and DNA. Comfort with IMFs unlocks the rest of this chapter and much of biochemistry.
The college version
Core Concepts
Dispersion (London) forces: the universal attraction
Every molecule and atom — even a single helium atom — experiences dispersion forces. Electrons move constantly, so at any instant the electron cloud of an atom may be slightly lopsided. That momentary temporary dipole polarizes a neighbor's electron cloud, inducing a matching dipole, and the two attract. The attraction flickers on and off, but on average it is real.
Two factors control dispersion strength:
- polarizability How easily an electron cloud distorts in response to charge Full entry → — how easily an electron cloud distorts. Bigger atoms with more, loosely held electrons are more polarizable.
- Surface contact — larger molecules can touch neighbors at more points, multiplying the attraction.
Among similar molecules, dispersion forces grow with molar mass. Among isomers with the same formula, the more elongated molecule has more surface contact than a compact, branched one, so it has stronger dispersion forces and a higher boiling point.
Dipole–dipole forces
When a molecule has permanent bond dipoles that do not cancel (a polar molecule like HCl or acetone), its positive end attracts its neighbor's negative end. These dipole–dipole forces are stronger than dispersion forces for small molecules of similar size, and they add on top of dispersion forces. They explain why polar molecules of similar molar mass boil higher than nonpolar ones: compare acetone (58 g/mol, bp 56 °C) with butane (58 g/mol, bp −0.5 °C).
Hydrogen bonding: a special strong dipole–dipole force
Hydrogen bonding is a particularly strong dipole–dipole attraction that forms when hydrogen is bonded to a highly electronegative atom — F, O, or N — and is attracted to a lone pair on another F, O, or N. The small, electron-poor hydrogen acts almost like a bare proton with a concentrated positive charge, so it pulls unusually hard on a neighbor's lone pair.
Hydrogen bonds are about ten times weaker than covalent bonds but far stronger than ordinary dipole–dipole forces. They cause water's anomalously high boiling point, ice floating, DNA base pairing, and protein folding. A hydrogen bond Strong dipole–dipole attraction between H bonded to N/O/F and a lone pair on N/O/F Full entry → needs a donor (H–F/O/N group) and an acceptor (a lone pair on F/O/N).
Ion–dipole forces
When an ion meets a polar molecule, the ion's charge strongly attracts the molecule's partial charges: water's positive end points toward a negative ion and vice versa. Ion–dipole forces are the strongest of the intermolecular interactions and explain why ionic solids like NaCl dissolve in water: the attraction between Na⁺ or Cl⁻ and water molecules overcomes the ionic bonds in the crystal.
Relative strength
Ordered from weakest to strongest:
dispersion < dipole–dipole < hydrogen bonding < ion–dipole
with the caveat that dispersion forces grow with size and can outrank dipole–dipole forces for large molecules. The total attraction — the net intermolecular force — decides the properties explored next: vapor pressure, boiling point, surface tension, viscosity, and whether a substance is a gas, liquid, or solid at room temperature.
How It Works / Step-by-Step Process
To predict which of two substances has the stronger IMFs and the higher boiling point:
- Identify the forces each substance can experience (dispersion always; dipole–dipole if polar; hydrogen bonding if H–N/O/F is present).
- Compare molar masses and shapes to judge dispersion strength.
- Weigh the strongest force present in each substance.
- Predict: more IMF strength → more energy to separate molecules → higher boiling point.
Common Confusions
| Common Confusion | Correct Understanding |
|---|---|
| "Hydrogen bonds hold the hydrogen atoms inside a water molecule." | No — the O–H bond inside water is covalent. Hydrogen bonds are between molecules: one molecule's H to another's O. |
| "Polar molecules only have dipole–dipole forces." | Polar molecules also have dispersion forces; all molecules do. |
| "Dispersion forces only occur in nonpolar molecules." | They occur in all molecules; for nonpolar ones they are simply the only IMF. |
| "Bigger molar mass always means higher boiling point." | Usually true among similar molecules, but strong hydrogen bonding (water vs H₂S) can reverse the trend. |
| "Hydrogen bonding requires any hydrogen atom." | It requires H bonded to N, O, or F, attracted to a lone pair on N, O, or F. C–H hydrogens do not hydrogen-bond. |
| "Ion–dipole forces are weaker than hydrogen bonds." | They are stronger — the full charge of an ion beats the partial charges of a dipole. |
Eli explains
The same idea, in plain words
Explain it like I’m 10
Molecules are like sticky magnets. Some are barely sticky at all (helium), some are medium-sticky (gasoline), and water molecules are extra sticky because their shape makes a special super-sticky patch called a hydrogen bond. How hard molecules stick together decides whether a substance is a gas that floats away, a liquid that pours, or a solid that holds its shape.
Worked example
Example 1: Why does water boil 262 °C higher than hydrogen sulfide?
H₂S has a higher molar mass (34.08 g/mol) than water (18.02 g/mol), so its dispersion forces are stronger. Yet water boils at 100 °C while H₂S boils at about −60 °C. The reason is hydrogen bonding:
- In water, each O–H bond places hydrogen next to the very electronegative oxygen, and oxygen carries two lone pairs. Water molecules form strong hydrogen bonds with each other.
- In H₂S, sulfur is less electronegative, and the S–H bond is not polar enough to support hydrogen bonding. H₂S molecules are held together only by dispersion and weak dipole–dipole forces.
Extra energy must break water's hydrogen bonds before molecules can escape as vapor, so water boils far higher despite its lower molar mass. Across the series H₂O, H₂S, H₂Se, H₂Te, water is the lightest yet boils the highest.
Example 2: Comparing boiling points of hydrocarbons
Rank these by predicted boiling point: butane (C₄H₁₀, 58 g/mol, straight chain), isobutane (C₄H₁₀, 58 g/mol, branched), and pentane (C₅H₁₂, 72 g/mol, straight chain).
Step 1 — forces present: all three are nonpolar, so only dispersion forces act.
Step 2 — molar mass: pentane (72 g/mol) has more electrons and stronger dispersion than the C₄ pair.
Step 3 — shape: butane is a straight chain with more surface contact; isobutane is branched and compact, so its molecules touch fewer neighbors.
Step 4 — prediction: pentane > butane > isobutane.
Measured values confirm: pentane 36 °C, butane −0.5 °C, isobutane −12 °C. Molar mass dominates first; among isomers, shape decides.
Key takeaways
- IMFs are attractions between molecules; bonds hold atoms within a molecule.
- Dispersion forces exist in all molecules and atoms; they increase with polarizability and surface contact (roughly, with molar mass).
- Dispersion < dipole–dipole < hydrogen bonding < ion–dipole (for molecules of comparable size).
- Hydrogen bonding requires H bonded to N, O, or F and a lone pair on N, O, or F.
- Stronger IMFs → higher boiling point, lower vapor pressure, higher surface tension, higher viscosity.
- More surface contact (less branching) → stronger dispersion → higher boiling point among isomers.
- Water's unusual properties (high boiling point, ice floating, strong cohesion) come from hydrogen bonding.
Check yourself
5 review questions from the chapter. Try each one, then open the answer.
List the four types of intermolecular forces in order of increasing strength.
Show answer
Dispersion < dipole–dipole < hydrogen bonding < ion–dipole (for molecules of comparable size).
Why do larger molecules generally have stronger dispersion forces?
Show answer
Larger molecules have more electrons, which are more easily polarized (higher polarizability), and more surface area for contact with neighbors — both strengthen dispersion forces.
What structural requirement must be met for hydrogen bonding to occur?
Show answer
Hydrogen must be covalently bonded to N, O, or F, and the molecule must have access to a lone pair on another N, O, or F (the acceptor).
Two isomers, one straight-chain and one branched, have the same formula. Which boils higher and why?
Show answer
The straight-chain isomer boils higher: it has more surface contact, so its dispersion forces are stronger than the compact branched form's.
Explain why NaCl, an ionic solid, dissolves in water despite strong ionic bonds.
Show answer
The ion–dipole attraction between Na⁺/Cl⁻ and polar water molecules is strong enough to pull the ions apart from the crystal and keep them hydrated in solution.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- intermolecular force (IMF)
- Attraction between separate molecules (or atoms)
- dispersion force
- Weak, momentary attraction from flickering electron clouds
- polarizability
- How easily an electron cloud distorts in response to charge
- dipole
- A molecule (or bond) with separated partial positive and negative ends
- dipole–dipole force
- Attraction between the positive end of one polar molecule and the negative end of another
- hydrogen bond
- Strong dipole–dipole attraction between H bonded to N/O/F and a lone pair on N/O/F
- ion–dipole force
- Attraction between an ion and a polar molecule
Sources & references
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