Chemistry: Atoms First 2e · Liquids and Solids
Intermolecular Forces
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In 30 seconds
Intermolecular forces (IMFs) are the attractions between molecules — as opposed to the covalent or ionic bonds within them. Weak (roughly 1–40 kJ/mol), they still decide whether a substance is a gas, liquid, or solid at room temperature, how high it boils, and whether it dissolves in water.
There are four main types, in order of increasing typical strength: dispersion (London) forces, dipole Separation of partial positive and negative charge in a molecule Full entry →–dipole forces, hydrogen bonding, and ion–dipole forces — all electrostatic attractions between opposite charges. This topic explains where each comes from, why dispersion is universal, and how IMF strength shows up in boiling point and enthalpy of vaporization Heat to vaporize one mole of liquid Full entry →.
Why this matters
IMFs explain the world around you: water's anomalously high boiling point (100 °C vs. −60 °C for H₂S) keeps it liquid and makes it the solvent of life; hydrogen bonding holds DNA's strands together and folds proteins. Practically, IMFs decide which solvents dissolve which drugs, why oil and water separate, how detergents work, and why some fuels are volatile fire hazards.
The college version
Core Concepts
The hierarchy of intermolecular forces
| Force | Typical energy (kJ/mol) | Occurs between | Origin |
|---|---|---|---|
| Dispersion (London) | 1–10 | all molecules | Fluctuating electron clouds create instantaneous dipoles |
| Dipole–dipole | 5–25 | polar molecules | Permanent partial charges attract |
| Hydrogen bond | 10–40 | H on N/O/F + lone pair on N/O/F | Strong dipole–dipole with a nearly bare proton |
| Ion–dipole | 40–600 | ions + polar molecules | Full ion charge attracts a molecular dipole |
A substance's total attraction is the sum of all types present.
Dispersion (London) forces — universal and often dominant
Electrons move constantly, so even a nonpolar molecule momentarily has more electron density on one side — a fleeting instantaneous dipole that polarizes a neighbor into an induced dipole. Brief though each fluctuation is, the time-averaged attraction is real.
Dispersion strength grows with polarizability How easily a molecule's electron cloud distorts Full entry → — how easily a molecule's electron cloud distorts. Bigger atoms and molecules, with more electrons held farther from the nucleus, are more polarizable, which is why boiling points climb down a group: F₂ (−188 °C), Cl₂ (−34 °C), Br₂ (59 °C), I₂ (184 °C), all held together by dispersion alone.
Dipole–dipole forces
Polar molecules have permanent dipoles from unequal electron sharing; the positive end of one attracts the negative end of a neighbor. Weaker than ionic attractions but additive — e.g., HCl is heavier than HF, yet HCl boils at −85 °C while HF boils at 20 °C, because HF hydrogen-bonds and HCl cannot.
Hydrogen bonding
A hydrogen bond Strong dipole–dipole attraction between H on N/O/F and a lone pair on N/O/F Full entry → is a strong dipole–dipole interaction between a hydrogen bonded to a highly electronegative atom (N, O, or F) and a lone pair on another N, O, or F. The H is left almost bare — a concentrated positive charge — so it grips a lone pair tightly. Hydrogen bonds (10–40 kJ/mol) are the strongest dipole-based IMF and are directional, organizing water's structure and DNA's double helix.
Ion–dipole forces
When an ion meets a polar molecule, the full ionic charge strongly attracts its dipole — the force that dissolves salt in water: Na⁺ is surrounded by water with oxygen ends inward, Cl⁻ by water with hydrogen ends inward. Ion–dipole is the strongest IMF type, yet far weaker than the ionic bonds inside the salt crystal.
From IMFs to bulk properties
Boiling point and enthalpy of vaporization rise as IMFs strengthen, since separating molecules costs more energy; at a given temperature, stronger IMFs also mean lower vapor pressure — the link to the next topic.
How It Works / Step-by-Step Process
To predict relative boiling points:
- Identify all IMF types present (dispersion always; dipole–dipole if polar; hydrogen bonding if H on N/O/F).
- Compare polarities: hydrogen bonding outranks dipole–dipole, which outranks dispersion only.
- If types tie, compare size and shape — more electrons and surface area mean stronger dispersion.
- Stronger total IMFs → higher boiling point, higher enthalpy of vaporization, lower vapor pressure.
Common Confusions
| Common Confusion | Correct Understanding |
|---|---|
| A hydrogen bond is a real chemical bond. | It is an intermolecular attraction, not a shared-electron bond — much weaker (10–40 kJ/mol) than covalent bonds. |
| Dispersion forces only occur in nonpolar molecules. | They occur between all molecules; they're just the only force for nonpolar ones. |
| Bigger molecules always have stronger IMFs. | Size matters only when IMF types tie; hydrogen bonding beats dispersion (HF boils higher than heavier HCl). |
| Any H atom can hydrogen-bond. | The H must be bonded to N, O, or F and approach a lone pair on another N, O, or F. |
| Dissolving salt breaks ionic bonds directly. | Ions are separated by ion–dipole stabilization with water, which repays the energy cost. |
| Higher boiling point means stronger covalent bonds. | Boiling point reflects intermolecular forces; covalent bonds survive boiling intact. |
Eli explains
The same idea, in plain words
Explain it like I’m 10
Molecules are like tiny magnets that pull on each other. Some pulls are weak and happen between every molecule, even ones with no plus or minus ends. Some are strong, like when a hydrogen atom grabs onto a nearby oxygen or nitrogen. Water boils much hotter than it "should" because its molecules hold onto each other with those strong grabs.
Worked example
Example 1: Boiling points of the halogens — dispersion only
F₂, Cl₂, Br₂, and I₂ are all nonpolar diatomic molecules; dispersion is the only IMF. As molar mass and electron count increase down the group, polarizability increases, so dispersion strengthens:
| Molecule | Molar mass (g/mol) | Boiling point (°C) |
|---|---|---|
| F₂ | 38 | −188 |
| Cl₂ | 71 | −34 |
| Br₂ | 160 | 59 |
| I₂ | 254 | 184 |
The trend is clean and monotonic — roughly twice the electrons roughly doubles the boiling point — classic evidence that dispersion alone can dominate phase behavior.
Example 2: Ethanol vs. dimethyl ether — same formula, different IMFs
Ethanol (CH₃CH₂OH) and dimethyl ether (CH₃OCH₃) both have formula C₂H₆O (molar mass 46 g/mol), so their dispersion forces are nearly identical. The difference is the O–H group: ethanol has one; dimethyl ether does not.
Ethanol can donate a hydrogen bond (its O–H) and accept one (oxygen lone pairs), so ethanol molecules hydrogen-bond to each other. Dimethyl ether has lone pairs but no O–H, so it can only accept hydrogen bonds — it cannot form them with itself, leaving dipole–dipole only.
Consequence: ethanol boils at 78 °C, dimethyl ether at −24 °C — a 100-degree gap from one O–H group. The same reasoning explains why water (two O–H groups) boils at 100 °C while heavier H₂S boils at −60 °C.
Key takeaways
- IMFs act between molecules; bonds act within them. IMFs are always much weaker.
- Dispersion forces operate between all molecules and often dominate for large molecules.
- Dispersion strength increases with polarizability: more electrons, bigger atoms, more surface area.
- Dipole–dipole requires polar molecules; hydrogen bonding requires H on N/O/F plus an N/O/F lone-pair acceptor.
- Hydrogen bonds are the strongest dipole-based IMF (10–40 kJ/mol) and are directional.
- Ion–dipole forces explain dissolving ionic compounds in polar solvents like water.
- Boiling point and enthalpy of vaporization rise with total IMF strength; vapor pressure falls.
- Compare molecules of similar size by IMF type; compare molecules of similar polarity by size (dispersion).
Check yourself
6 review questions from the chapter. Try each one, then open the answer.
List the four types of intermolecular forces in order of increasing typical strength.
Show answer
Dispersion (London), dipole–dipole, hydrogen bonding, ion–dipole.
Why do all molecules experience dispersion forces, even nonpolar ones?
Show answer
Electrons move constantly, so every molecule briefly develops an instantaneous dipole that induces one in its neighbor — the attraction is universal.
Explain why Br₂ is a liquid at room temperature while F₂ is a gas.
Show answer
Both are nonpolar with only dispersion forces; Br₂ has more electrons (more polarizable), so stronger dispersion holds it liquid at room temperature.
Why does HF boil at 20 °C even though HCl is heavier and boils at −85 °C?
Show answer
HF can form hydrogen bonds; HCl cannot (Cl is too weakly electronegative), so it relies on weaker dipole–dipole and dispersion.
What two structural requirements must be met for hydrogen bonding?
Show answer
A hydrogen bonded to N, O, or F, plus a lone pair on another N, O, or F to accept it.
Which IMF type explains why NaCl dissolves in water, and how does it work?
Show answer
Ion–dipole forces: water's dipoles orient around each ion (oxygen toward Na⁺, hydrogens toward Cl⁻), stabilizing the separated ions.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- intermolecular force
- Attraction between separate molecules
- dispersion (London) force
- Attraction from fluctuating electron clouds creating temporary dipoles
- polarizability
- How easily a molecule's electron cloud distorts
- dipole
- Separation of partial positive and negative charge in a molecule
- dipole–dipole force
- Attraction between permanent dipoles of polar molecules
- hydrogen bond
- Strong dipole–dipole attraction between H on N/O/F and a lone pair on N/O/F
- ion–dipole force
- Attraction between an ion and a polar molecule's dipole
- enthalpy of vaporization
- Heat to vaporize one mole of liquid
Sources & references
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