If organic chemistry has a native language, it is the curved arrow, and this is the single most important skill to master early because every mechanism you will ever draw is built from it. The core idea is simple and rigid: a curved arrow shows the movement of electrons, never the movement of atoms. Get this backward and nothing downstream will make sense.

Each arrow has two ends that mean specific things. The tail is placed exactly on the source of the electrons — either a lone pair or a bond (a pair of bonding electrons). The head points to exactly where those electrons are going — toward an atom, or into the space between two atoms to form a new bond. There are two kinds of arrows: a double-barbed (full) arrow moves an entire pair of electrons and is what you use for the vast majority of reactions (these are "polar" or "ionic" mechanisms). A single-barbed arrow, drawn like a fishhook, moves only one electron and is reserved specifically for radical mechanisms. Do not mix them up — using a fishhook in a polar mechanism signals a fundamentally different (and usually wrong) process.

Two rules are absolute. First, only electrons move; the atomic nuclei stay put. Second, electrons are conserved — you can never create or destroy them with an arrow, only relocate them. A helpful way to self-check any arrow: an atom that receives a new bond must be able to accommodate it (carbon that already has four bonds cannot simply accept a fifth without simultaneously losing one), and an atom that donates electrons develops a more positive character while the recipient develops more negative character.

With arrows in hand, we can properly understand resonance. Resonance is what you invoke when a single Lewis structure honestly cannot describe a molecule's real electron distribution. The classic example is the carboxylate ion (a carboxylic acid that has lost its proton, –COO⁻). You can draw it two ways: the double bond on the left oxygen with the negative charge on the right, or the double bond on the right oxygen with the negative charge on the left. Experiment shows neither is correct — the two C–O bonds are identical in length, halfway between a single and double bond, and the negative charge is spread equally over both oxygens.

Here is the concept students most often get wrong, so it deserves emphasis: the molecule does not flip back and forth between the two structures. It is not sometimes-one, sometimes-the-other. The real molecule is a single, unchanging structure called the resonance hybrid, which is a blend — a weighted average — of the contributing structures we are able to draw. We draw multiple structures only because our simple line-and-dot notation is too crude to depict "half a bond" or "half a charge" in one picture. The individual resonance structures are the limitation of our drawing system, not real states the molecule occupies. (The double-headed arrow ↔ between resonance structures means "these are two depictions of one thing," and is completely different from the equilibrium arrows ⇌, which mean two real species interconverting. Never confuse the two.)

Not all resonance structures contribute equally. To judge which ones weight the true hybrid most heavily — which is the same as judging where the electrons really prefer to be — use these criteria:

Structures with complete octets on all atoms contribute most.
Structures with fewer formal charges and less charge separation contribute more.
Structures that place any negative charge on the most electronegative atom (and positive charge on the least electronegative atom) contribute more, because electronegative atoms are comfortable holding negative charge.

The payoff of resonance is stabilization. Spreading electrons (and charge) over several atoms rather than concentrating them on one atom always lowers energy — a delocalized charge is a more stable charge. The more good resonance structures you can draw, the more stabilized the species. This one principle quietly explains an enormous amount of what is to come: why carboxylic acids are far more acidic than alcohols, why certain carbocations form preferentially, why amides barely act as bases, and why benzene is extraordinarily unreactive for something with three "double bonds." When later material seems to hinge on stability, resonance is very often the reason.

"Push the Arrows"