heterogeneous equilibrium
Picture a closed jar holding chunks of solid limestone, with carbon dioxide gas above and the solid slowly breaking down and re-forming. The reaction involves more than one phase at once — solid and gas — sharing the same container. An equilibrium that spans two or more distinct phases like this is called a heterogeneous equilibrium, in contrast to a homogeneous one where everything is in a single phase.
The defining feature shows up in the equilibrium constant: pure solids and pure liquids are left out of the expression entirely. The reason is that their 'concentration' — really their activity — does not change as the reaction proceeds; a block of solid has the same density whether it is large or small. So only the gases and dissolved species appear in K. Adding or removing some pure solid, as long as a little remains, does not shift the equilibrium at all.
This is genuinely useful and a frequent source of confusion. Students often try to put a solid into the equilibrium expression and get the wrong answer. The rule is firm: include gases (as pressures or concentrations) and dissolved solutes, but never a pure solid or pure liquid solvent. The honest refinement is that this works because the activity of a pure condensed phase is defined to be 1.
For CaCO3(s) ⇌ CaO(s) + CO2(g), both calcium carbonate and calcium oxide are pure solids, so they drop out. The equilibrium expression is simply Kp = p(CO2): the whole equilibrium is captured by the pressure of one gas.
Pure solids and liquids are left out of the equilibrium expression.
Leaving a pure solid out does not mean it is irrelevant — at least a trace must be present for the equilibrium to exist. If the solid runs out completely, there is simply no equilibrium of that kind to speak of.