This topic explains the physical behavior of molecules — why some are gases and some are waxy solids at room temperature, why some dissolve in water and some float on top of it — and it all comes down to the forces between separate molecules. It is essential to keep two very different kinds of forces straight. Intramolecular forces are the covalent bonds within a single molecule, holding its atoms together — these are strong. Intermolecular forces (IMFs) are the attractions between one molecule and its neighbors — these are much weaker, but they are what you must overcome to melt or boil a substance, and they determine solubility. When water boils, you are not breaking the strong O–H bonds inside each water molecule; you are only pulling water molecules away from each other, breaking the weaker intermolecular attractions.

There are three main types of IMF, and it is worth understanding the physical origin of each rather than just memorizing the ranking. From weakest to strongest:

London dispersion forces exist in every molecule, including completely nonpolar ones. Their origin is beautifully simple: electrons are in constant motion, so at any given instant they may be lopsided within a molecule, creating a fleeting, temporary dipole. That momentary dipole induces an opposite dipole in a neighboring molecule, and the two attract. Any single such attraction is tiny, but they add up. They get stronger with more surface area and more electrons (bigger, longer molecules are more "polarizable" — their electron clouds are squishier and easier to distort). This is why boiling point climbs steadily as you lengthen a carbon chain, and why a branched molecule boils lower than its straight-chain isomer: branching makes the molecule more compact and spherical, reducing the surface area available for neighbors to touch.
Dipole–dipole forces occur in polar molecules — those with a permanent, built-in dipole because of unevenly shared electrons. The permanently positive end (δ+) of one molecule is attracted to the permanently negative end (δ−) of another. These are stronger than dispersion forces of comparable size.
Hydrogen bonding is the strongest ordinary intermolecular force and, despite its name, is not an actual bond — it is a specially strong dipole–dipole attraction. It occurs under one specific condition: a hydrogen atom must be bonded directly to N, O, or F (small, highly electronegative atoms), which leaves that hydrogen very electron-poor and δ+; it is then strongly attracted to a lone pair on an N, O, or F of a neighboring molecule. Hydrogen bonding dominates the properties of water, alcohols, amines, and carboxylic acids, and it is the main reason such molecules have unexpectedly high boiling points and dissolve well in water.

Whether a molecule is polar at all depends on two things working together: the polarity of its individual bonds, and the molecule's overall geometry. This trips people up: a molecule can contain strongly polar bonds yet be nonpolar overall if its shape is symmetric enough that the individual bond dipoles point in opposing directions and cancel out. Carbon dioxide (O=C=O) is the textbook case — each C=O bond is polar, but the molecule is linear and symmetric, so the two dipoles pull equally in opposite directions and the net dipole is zero. So always check both the bonds and the shape before declaring a molecule polar or nonpolar.

All of this culminates in the rule that governs solubility: "like dissolves like." Polar (and hydrogen-bonding) substances dissolve in polar solvents such as water, because the solute and solvent can form the same kinds of favorable attractions with each other that they formed with themselves. Nonpolar substances dissolve in nonpolar solvents for the same reason. Mixing unlike substances — say, oil (nonpolar) and water (polar) — is unfavorable because the strongly hydrogen-bonding water molecules would rather cling to each other than surround a nonpolar oil molecule, so the two separate into layers.

For organic molecules specifically, solubility becomes a tug-of-war within a single molecule. A small polar or hydrogen-bonding functional group like –OH (hydroxyl) or –COOH (carboxyl) pulls the molecule toward water-solubility. But every molecule also has a nonpolar hydrocarbon portion (its carbon–hydrogen skeleton) that pulls the other way, toward water-insolubility. In a small molecule like methanol or ethanol, the polar –OH dominates and the molecule mixes freely with water. But as the nonpolar carbon chain grows longer, it eventually overwhelms the little polar group, and the molecule stops dissolving in water — a long-chain alcohol behaves much more like an oil than like methanol. Learning to weigh a molecule's polar "handle" against the size of its nonpolar "body" lets you predict solubility on sight, a skill you will lean on when you reach extractions and purifications later in the course.

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