$word_{count}$: 220

Ksp governs dissolution equilibria of sparingly soluble salts. For AxBy: Ksp = [A]^x[B]^y = $x^x$ * $y^y$ * s^(x+y). Common types: AB (Ksp = $s^2$), AB2 (Ksp = 4$s^3$), A2B (Ksp = 4$s^3$), AB3 (Ksp = 27$s^4$), A2B3 (Ksp = 108$s^5$). Precipitation occurs when ionic product IP > Ksp. The common ion effect reduces solubility dramatically by adding an ion already present in the equilibrium. When comparing solubility of salts with different stoichiometries, you must calculate s explicitly — direct Ksp comparison is only valid for the same stoichiometric type. Selective precipitation exploits Ksp differences: the less soluble salt (lower Ksp for same type, or lower [reagent] needed) precipitates first. Solubility increases in acidic conditions for salts of weak acids (CaCO3, CaF2) because H+ consumes the anion. Salts of strong acids (AgCl) are unaffected by pH. Complex formation (e.g., Ag+ + 2NH3 -> Ag(NH3)2+) increases apparent solubility beyond Ksp prediction. Overall K for dissolution + complexation = Ksp * Kf.