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Ca +

Precipitate from Common Ions. Ksp = ] [ Ca + ] + [ OH¯ ]. Ca(OH) 2  Ca + + OH¯. A saturated solution starts out with the solid salt Ca(OH) 2 at equilibrium with a specific concentration of dissolved ions produced from that salt .

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Ca +

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  1. Precipitate from Common Ions Ksp = ][Ca+] + [OH¯] Ca(OH)2Ca+ + OH¯ A saturated solution starts out with the solid salt Ca(OH)2 at equilibrium with a specific concentration of dissolved ions produced from that salt. This is the normal solubility or concentration, in Molarity, of the dissociated Ca(OH)2. The Kspis the constant for the measure of equilibrium for solubility. Ca+ Ca+ OH¯ OH¯ OH¯ OH¯ Ca+ OH¯ OH¯ Ca+ Ca+ OH¯ OH¯ OH¯ OH¯ Ca(OH)2 Ca(OH)2

  2. Precipitate from Common Ions 5.02 x 10-6 = [Ca+] + [OH¯] Ca(OH)2Ca+ + OH¯ 34.0 = [Na+] + [OH¯] NaOHNa+ + OH¯ However, if there are ions from another source, in this case NaOH, that are dissociated according to a higher Ksp, then since the Ksp of the Ca(OH)2 is lower, some of the OH- ions have to join back with some Ca+ ions to turn back into solid Ca(OH)2 so they are not counted in the Ksp expression This is Le Chatelier’s Principle at work. Note: this is where the term “common ions” comes from – the NaOH has the OH- ion in common with the Ca(OH)2 Ca+ Ca+ OH¯ OH- OH- OH¯ Ca+ OH¯ OH¯ OH¯ OH¯ Na+ Na+ Ca+ Ca+ OH¯ OH¯ OH¯ OH¯ Ca(OH)2 Ca(OH)2

  3. Precipitate from Common Ions Ksp = ][Ca+] + [OH¯] Ca(OH)2Ca+ + OH¯ 34.0 = [Na+] + [OH¯] NaOHNa+ + OH¯ Some of the Ca2+ ions reassociate with the OH- ions and turn to solid (or precipitate out). Ca+ Ca(OH)2 OH¯ OH¯ Ca+ OH¯ OH¯ OH¯ OH¯ Na+ Na+ Ca+ Ca(OH)2 OH¯ OH¯ Ca(OH)2 Ca(OH)2

  4. Precipitate from Common Ions Ksp = ][Ca+] + [OH¯] Ca(OH)2Ca+ + OH¯ 34.0 = [Na+] + [OH¯] NaOHNa+ + OH¯ The solution is again at equilibrium even though the concentrations of Ca+ions and OH-ions from the Ca(OH)2 salt are lower than previous. Ca+ OH¯ OH¯ Ca+ OH¯ OH¯ OH¯ OH¯ Na+ Na+ Ca+ OH¯ OH¯ Ca(OH)2 Ca(OH)2 Ca(OH)2 Ca(OH)2

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