equilibrium constant
/ K (often 'K-ee' for Kc) /
Imagine you let many different reactions settle down, each in its own jar, all at the same temperature. For each reaction you could ask: once it stops changing, what is the recipe of the final mix — how much product compared to leftover reactant? The equilibrium constant is a single number that captures that final recipe for a given reaction at a given temperature. It tells you, in advance, roughly how far the reaction will go.
For a balanced reaction, the equilibrium constant K is the ratio of products to reactants at equilibrium, with each substance raised to the power of its coefficient in the equation. A large K (much greater than 1) means the mix ends up rich in products — the reaction goes nearly to completion. A small K (much less than 1) means it ends up rich in reactants — barely anything reacts. A K near 1 means a genuine mixture of both. Importantly, K depends only on temperature, not on how much you started with or how fast you got there.
The honest subtlety is that the things you put into the ratio should really be activities (effective concentrations) or, for gases, fugacities, which is why the truly rigorous version is dimensionless. In everyday school work we approximate with concentrations (Kc) or partial pressures (Kp). And K tells you the destination, not the journey: a reaction can have a huge K yet sit untouched for years if it has no fast pathway.
For the reaction H2 + I2 ⇌ 2HI at 458 °C, Kc is about 49. Since 49 is comfortably above 1, an equilibrium mixture is mostly hydrogen iodide, with only modest leftover hydrogen and iodine — the reaction strongly favours products at that temperature.
K above 1 favours products; K below 1 favours reactants.
K is fixed by temperature alone. Changing concentrations or pressure shifts the position of equilibrium but leaves K unchanged; only a change in temperature changes K itself.