A tug-of-war over electrons
You already know that a covalent bond is two atoms sharing a pair of electrons. But "sharing" almost never means "sharing equally." Every atom has its own appetite for electrons, called its [[org-electronegativity|electronegativity]] — think of it as how hard the nucleus pulls on the electrons in a bond. Fluorine and oxygen are greedy; carbon and hydrogen are mild; metals are generous. When two atoms of different appetite bond, the shared pair drifts toward the greedier one, like a rope in a tug-of-war sliding toward the stronger team.
When the pull is uneven, the bond becomes a [[org-bond-polarity|polar bond]]. The greedier atom ends up with slightly more than its fair share of electron density, so it carries a small negative charge, written delta-minus. Its partner, left a little short, carries an equal delta-plus. These are partial charges — not full ions, just a faint, permanent imbalance. In an O-H bond the oxygen is delta-minus and the hydrogen delta-plus; in a C=O bond the oxygen pulls and the carbon is left delta-plus. Hold onto that last one — it is the single most important polar bond in all of organic chemistry.
From polar bonds to a polar molecule
A single polar bond has a direction and a size — together they make a [[org-dipole-moment|dipole]], a little arrow pointing from delta-plus to delta-minus. But a molecule has many bonds, and here the shapes you learned last guide come roaring back. Each bond dipole is a vector, and the molecule's overall polarity is their sum. If the arrows cancel, the molecule is nonpolar even though its bonds are polar; if they reinforce, the molecule has a strong net dipole.
Carbon dioxide is the classic trap. O=C=O has two strongly polar C=O bonds, yet the molecule is linear, so the two dipole arrows point dead opposite and cancel to zero — CO2 is nonpolar. Water draws the same two O-H dipoles but its bent shape (the lone pairs push the H's down) leaves the arrows pointing the same general way, so they add up and water is strongly polar. Same kind of bonds, opposite verdict, decided entirely by geometry. Always ask two questions: are the bonds polar, and does the shape let their dipoles survive?
O=C=O linear <-O C O-> dipoles cancel -> nonpolar
H H bent both arrows tilt the same way
\ / -> dipoles ADD -> polar
O(:)(:)Polarity is a map of where reactions happen
Partial charges are not just trivia about boiling points — they are a treasure map for reactivity. Opposite charges attract, so electron-rich species seek out the delta-plus spots and electron-poor species seek out the delta-minus spots. A [[nucleophile|nucleophile]] ("nucleus-lover," electron-rich) goes hunting for a delta-plus carbon; an [[electrophile|electrophile]] ("electron-lover," electron-poor) goes hunting for a delta-minus, electron-rich site. Long before you know any named reaction, you can often predict where a molecule will get attacked just by shading in its partial charges.
Take the carbonyl C=O again. The carbon is delta-plus, sitting there like a target, which is exactly why nucleophiles add to aldehydes and ketones at the carbon. In a haloalkane like CH3CH2Br, the carbon attached to bromine is delta-plus because bromine pulled electron density away, marking it as the spot a nucleophile will attack. The polarity tells you the address of the reaction; the mechanism — which you will meet in later rungs — tells you how the delivery happens. For now, just train the habit: see a polar bond, picture the partial charges, and you have a head start on the chemistry.
The three forces between molecules
So far we have looked inside a molecule. Now zoom out to a whole flask of them. The bonds holding atoms together are strong (covalent). The far weaker attractions that hold separate molecules near each other are the [[intermolecular-forces|intermolecular forces]], often called van der Waals forces. They are the reason anything condenses into a liquid or freezes into a solid at all. Three of them matter for organic molecules, and they come in increasing strength.
- London dispersion forces — the universal one, present in everything, even nonpolar molecules. Electrons jitter; for an instant a molecule has a fleeting delta-plus side and delta-minus side, and it nudges its neighbor's electrons to do the same. These flickering, in-sync dipoles attract. Weak per contact, but they grow with surface area: a long, floppy alkane chain has far more touching surface than a small round one, so bigger molecules cling harder.
- Dipole-dipole forces — for molecules with a permanent net dipole. The delta-plus end of one molecule lines up with the delta-minus end of its neighbor, like tiny bar magnets snapping head-to-tail. Stronger than dispersion at the same molecular size, so a polar molecule generally boils higher than a nonpolar one of similar weight.
- Hydrogen bonding — the strongest of the three, and a special, supercharged dipole-dipole. It needs an H bonded directly to N, O, or F (where the H is left very delta-plus and almost bare), reaching toward a lone pair on an N, O, or F of another molecule. This is why water, alcohols, and amides cling so fiercely. It is still far weaker than a real covalent bond — do not confuse the two.
Why this sets boiling points and what dissolves what
Boiling means tearing molecules apart from each other into a gas, so the stronger the intermolecular forces, the more energy it takes, and the higher the boiling point. Compare three molecules of similar weight: ethane (CH3CH3, only dispersion) boils around -89 C; the slightly polar fluoroethane is higher; and ethanol (CH3CH2OH, which can hydrogen bond) boils at +78 C. Same neighborhood of size, wildly different boiling points — that whole spread is the intermolecular forces talking.
The same logic governs solubility, captured in the old chemist's mantra "like dissolves like." A solvent dissolves a solute when the forces between them are about as friendly as the forces each had on its own. Water, knit together by hydrogen bonds, happily dissolves sugar and salt and ethanol — partners that can hydrogen bond back. But it shuns oil: a long hydrocarbon offers only dispersion forces and cannot pay water's hydrogen-bonding entry fee, so water would rather hold onto itself and the oil clumps apart. Polar dissolves polar; nonpolar dissolves nonpolar.
There is even a tug-of-war within a single molecule. Ethanol's tiny O-H end is polar and water-loving; ethanol dissolves freely. Stretch that chain out to an eight-carbon alcohol and the long, greasy hydrocarbon tail starts to dominate — solubility in water collapses. So polarity, dipoles, and intermolecular forces are not separate facts to memorize; they are one connected story, running from a single electronegativity difference all the way up to whether your compound is a gas, a puddle, or a powder, and whether it will dissolve in water or float on it.