Covalent bonds hold individual molecules together, but they do not decide if a substance is a solid, liquid, or gas at room temperature. That role belongs to the electrostatic forces of attraction between neighboring molecules. By the end of this page, you'll be able to identify the three types of intermolecular forces, rank substances by their strength, and predict bulk physical properties like boiling point, vapor pressure, and solubility.
To understand molecular behavior, we must distinguish between two very different electrical forces. The strong covalent bonds holding the hydrogen and oxygen atoms together inside a single water molecule are intramolecular bonds. These require massive chemical energy to break.
In contrast, the relatively weak electrical attractions that pull separate, neighboring water molecules toward one another are intermolecular forces (commonly abbreviated as IMFs). These are the forces you overcome during physical phase changes.
When you boil a kettle of water, the steam rising from the spout is still H2O. You have not broken the covalent bonds to separate hydrogen and oxygen gas. You have only added enough thermal energy to break the weak IMFs, letting individual water molecules escape from each other into the air.
Covalent substances experience three primary types of intermolecular forces. They are all electrostatic (opposite charges attracting), but they differ in how their charges are created and how long they last.
Temporary, induced dipoles. Electrons are in constant motion. By chance, they occasionally bunch up on one side of a molecule, creating a brief, temporary partial negative charge (δ−) and leaving the other side positive (δ+). This temporary dipole repels electrons in a neighboring molecule, inducing a matching dipole. The resulting weak attraction is a London dispersion force (or LDF). LDF is present in all molecules. Larger electron clouds (heavier atoms or more atoms) shift more easily, leading to stronger LDF.
Permanent polar attractions. If two nonmetals have mismatched electronegativities, they share electrons unequally. This permanent polar bond creates permanent partial charges (δ+ and δ−). When polar molecules get close, they align so their opposite poles attract. This permanent electrostatic pull is a dipole-dipole force.
Super-charged polar attraction. When hydrogen is bonded directly to nitrogen, oxygen, or fluorine (the three most electronegative elements, which also carry highly concentrated lone pairs), the electronegativity difference is massive. The electronegative atom pulls the shared electron density almost entirely away, leaving hydrogen's single proton completely unshielded. This bare, highly positive hydrogen attracts a negative lone pair on a neighboring molecule's N, O, or F. Despite the name, hydrogen bonding is not a true chemical bond; it is a very strong intermolecular force.
Select a substance. Slide the temperature slider from Cool (Solid) to Hot (Gas). Watch how the electrostatic attractions (dashed lines) respond to molecular movement while the chemical bonds (solid lines) remain unchanged.
Carbon dioxide is nonpolar and only attracts neighbors via weak London dispersion forces. It sublimes directly to gas at standard pressure because its attractions are easily overcome by minimal thermal energy.
Because we cannot see individual molecules, we must infer their electrostatic attraction by running macroscopic investigations. Two classic laboratory tests gather this evidence: evaporation cooling curves and surface tension beads.
When a liquid evaporates, it absorbs energy from its surroundings to break its IMFs. The faster it evaporates, the more rapidly it cools. A thermometer bulb wrapped in a soaked tissue will register a temperature drop. Liquids with weaker IMFs evaporate much faster, creating steeper temperature drops.
Those three molecules aren't three unrelated cases — they're three points on one continuous slide from weak to strong IMFs. Drag the slider and watch the same cooling curve shift.
Molecules on the surface of a liquid are pulled inward by attractions from neighboring molecules below them. This inward force is surface tension. When placed on a flat coin, water molecules pull on each other so strongly that they bead up into a high, rounded dome. Acetone, with weaker forces, spreads out flat and spills off the edge.
Consider three molecules of similar size: methane (CH₄, nonpolar), hydrogen chloride (HCl, polar), and ammonia (NH₃, polar with N–H bonds). Rank them in order of increasing boiling point.
CH₄ (LDF, weakest) < HCl (Dipole-Dipole, moderate) < NH₃ (Hydrogen bonding, strongest). Because methane is nonpolar, its molecules are attracted only by weak London dispersion forces. It boils at a very low −161°C. Polar HCl molecules align to pull dipole-to-dipole, boiling at −85°C. Ammonia's polar N–H bonds form strong hydrogen bonds, requiring the most thermal energy to break, leading to a boiling point of −33°C.
Hexane (C₆H₁₄) is a completely nonpolar solvent found in gasoline. If you drop a nonpolar oil stain into water (polar) and into hexane (nonpolar), where will the oil dissolve?
Oil only dissolves in hexane. The chemical rule of solubility is "like dissolves like." Polar water molecules are attracted to each other by strong hydrogen bonds, which squeeze nonpolar oil molecules out of the way (they cannot form favorable attractions with oil). Nonpolar hexane molecules interact with nonpolar oil molecules via LDF, letting them mix freely. This is why water alone cannot wash grease off your hands.
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Covalent bonds inside a molecule are . The attractions between separate molecules are . Temporary attractions from shifting electron clouds are . Attractions between permanently polar molecules are . The exceptionally strong IMF that forms when hydrogen bonds directly to N, O, or F is . Stronger IMFs hold molecules tighter, leading to a higher but a lower . The inward pull on a liquid's surface that makes water form beads is .
Explain why water (H₂O) has a much higher boiling point than carbon dioxide (CO₂), even though carbon dioxide is a much larger and heavier molecule.
Explain it in plain English. Link the types of forces to the energy needed to boil the liquid.
Water molecules can form hydrogen bonds (O–H bonds present), which are exceptionally strong intermolecular attractions. Carbon dioxide is a nonpolar molecule and can only form weak London dispersion forces (LDF). Even though CO₂ has a larger electron cloud, its weak LDF are much easier to overcome than water's strong hydrogen bonds, meaning water requires significantly more thermal energy (higher temperature) to boil.
Give yourself a point for each idea you actually wrote down. The flag (⚑) marks the step that separates a complete answer from a partial one.
Acetone (C₃H₆O) forms dipole-dipole attractions. Ethanol (C₂H₅OH) forms hydrogen bonds. Compare the relative strengths of their intermolecular forces, predict which substance will evaporate more rapidly at room temperature, and explain how this evaporation rate relates to their vapor pressures.
Self-score: 4 = all four · 3 = missing the connection to gas pressure · 2 = relative strengths + evaporation prediction, no molecular escape explanation · ≤1 = simple definition of hydrogen bonds/dipole-dipole only.
Geckos walk on glass using intermolecular forces. Geckos scale sheer vertical walls using millions of microscopic hairs (setae) on their feet. Because the hairs can get incredibly close to smooth glass, the gecko's grip relies entirely on the weak, collective attraction of London dispersion forces (LDFs).