Two ways to count a gas
So far we measured "how much" of each substance by concentration. But for gases there's a second, equally natural yardstick: partial pressure — the share of the total pressure that one gas contributes, as if it had the whole container to itself. A gas that is more crowded pushes harder, so its partial pressure rises and falls right alongside its concentration. Either yardstick faithfully tracks "how much."
Because we have two yardsticks, we get two flavours of equilibrium constant — together called Kc and Kp. *Kc* is the products-over-reactants fraction built from concentrations; *Kp* is the very same fraction built from partial pressures. They describe the same equilibrium, just in different currencies, the way a price can be quoted in dollars or in euros.
When solids and pure liquids quietly disappear
Now a subtlety that surprises every newcomer. Many reactions involve a mix of phases — a gas reacting with a chunk of solid, say, or a solid dissolving in a liquid. An equilibrium spanning more than one phase is called a heterogeneous equilibrium. And here is the odd rule: when you write *K* for such a reaction, pure solids and pure liquids are left out entirely.
Why would a chunk of solid not count? The intuition: "how crowded" is the wrong question for a pure solid. A lump of chalk has the same internal packing whether it's a pebble or a boulder — adding more chalk just makes a bigger lump, not a denser one. Since its "crowdedness" never changes, it contributes a constant that we simply fold into *K* and stop writing. The same goes for a pure liquid like the water in which something dissolves.
Measuring how far something falls apart
A whole family of equilibria are about a single substance splitting into pieces — a molecule breaking into smaller ones, or a salt separating into charged fragments. To describe how far that splitting goes, chemists use the degree of dissociation: the fraction of the original substance that has come apart at equilibrium. It runs from 0 (nothing split) to 1 (completely split), and is often quoted as a percentage.
Degree of dissociation and the equilibrium constant are two views of the same thing — if you know one, a little algebra hands you the other. A large *K* means the splitting is favourable, which shows up as a degree of dissociation close to 1. The lovely part is that the degree of dissociation often *changes* with conditions even though *K* stays fixed. Dilute a solution, for example, and weak substances tend to split more, nudging the degree upward without touching *K* at all.
Crowding the products: the common-ion effect
Here is a beautiful pay-off where Le Chatelier and dissociation meet. Suppose a salt is partly dissolved, sitting at equilibrium with the charged fragments it released. Now you stir in a second, different salt that happens to share one of those same fragments. You've just crowded the product side — and the equilibrium responds by shifting backward, pushing some of the dissolved salt back to solid. This is the common-ion effect.
It feels almost paradoxical at first: adding *more* dissolved material makes the original salt *less* soluble. But through the lens of concentration and *Q* it's perfectly natural — you raised the products, so *Q* shot above *K*, and the reaction ran backward to bring it home. This single effect explains why adding common salt can purify other salts, and why your body buffers itself against sudden changes in acidity.
Why all this fine print matters
None of these wrinkles is busywork. Leaving solids out of *K* is what lets us predict how a mineral dissolves underground. Choosing *Kp* over *Kc* is how engineers reason about gas reactors without juggling volumes. The degree of dissociation is the number a pharmacist needs to know how much of a drug is in its active, split-apart form. The fine print is where equilibrium meets the real, messy world.