Cornell Notes — Subtopic: Ksp and Common Ion Effect — Notes

Cue Column | Notes Column

Question / Keyword	Detail / Answer
Ksp definition for AmBn	$K_{sp} = [$A^{n+}$]^m[$B^{m-}$]^n$ — product of ion concentrations at saturation, each raised to its stoichiometric coefficient
Solubility s of AgCl in pure water?	AgCl ⇌ $Ag^{+}$ + $Cl^{-}$ ; Ksp = $s^{2}$ ; s = √Ksp = √( $1.8 \times 10^{-10}$ ) = $1.34 \times 10^{-5}$ M
Solubility of AgCl in 0.1 M NaCl?	[ $Cl^{-}$ ] ≈ 0.1 M (common ion); s = Ksp/0.1 = $1.8 \times 10^{-10}$ /0.1 = $1.8 \times 10^{-9}$ M — drastically reduced
Ksp of $Ag_{2}CrO_{4}$ expression	$Ag_{2}CrO_{4}$ ⇌ 2 $Ag^{+}$ + $CrO_{4}^{2-}$ ; Ksp = [ $Ag^{+}$ ]^{2}[ $CrO_{4}^{2-}$ ] = (2s)^{2}(s) = 4 $s^{3}$
Condition for precipitation	Ionic product Qsp > Ksp → precipitation occurs
Selective precipitation principle	Add precipitating agent slowly; the salt with the LOWER Ksp precipitates first
Effect of common ion on ionization	Suppresses ionization of weak electrolyte; e.g., HCl lowers ionization of acetic acid

Summary

The solubility product Ksp is an equilibrium constant for the dissolution of sparingly soluble salts. In pure water, solubility s is calculated from Ksp using stoichiometry. In the presence of a common ion, solubility is dramatically reduced because the equilibrium shifts backward (Le Chatelier). Selective precipitation exploits differences in Ksp values — the less soluble salt precipitates first. The common ion effect also suppresses weak-acid ionization, forming the basis of buffer chemistry.