equilibrium constant
Think of a single number that scores how a reaction prefers to end up — like a verdict that says 'mostly finished' or 'barely got started'. The equilibrium constant, written K, is that score. A big K means the reaction runs almost to completion and the flask ends up full of products; a tiny K means it hardly proceeds and you are left mostly with what you started with.
Concretely, K is built from the concentrations (strictly, the activities) of products and reactants once equilibrium is reached: you multiply the product concentrations together, each raised to its coefficient in the balanced equation, and divide by the reactant concentrations treated the same way. For a given reaction at a given temperature this ratio always lands on the same value, no matter how much of each substance you start with — that constancy is the whole point.
K is the backbone of quantitative analysis: it tells you how completely a titration reaction will go, how little of an ion will stay dissolved, how much of a coloured complex will form. The caveat to keep honest is temperature — K is a constant only at a fixed temperature. Heat or cool the system and the number shifts, which is why careful lab work controls temperature and why tabulated K values always come with a temperature attached.
For the dissociation of acetic acid, K (here called Ka) is about 1.8 x 10^-5 — a small number, telling you that in water only a tiny fraction of the acid molecules let go of their proton at any moment.
Small Ka = weak acid; most molecules stay intact.
K has many specialised faces with their own symbols — Ka for acids, Ksp for solubility, Kf for complex formation, Kw for water — but they are all the same idea: a fixed ratio of products to reactants at equilibrium. Pure solids and pure liquids are left out of the expression because their 'concentration' does not change.