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The reaction quotient Q represents the product of ion concentrations at any stage of a reaction, while Ksp is the product of ion concentrations specifically at equilibrium. Q changes as ions are added or removed, whereas Ksp remains constant for a given compound at a fixed temperature. Comparing Q to Ksp predicts whether precipitation will occur.

Calculate the reaction quotient Q from the ion concentrations in the mixed solution. If Q is less than Ksp, no precipitation occurs and the solution is unsaturated. If Q exceeds Ksp, the solution is supersaturated and precipitation will occur until Q equals Ksp and equilibrium is established.

Selective precipitation removes individual ions by gradually increasing the counter ion concentration. The metal ion forming the least soluble compound precipitates first at the lowest counter ion concentration. Other ions subsequently precipitate as their compound solubilities are reached, allowing separation based on differing Ksp values.

Lead(II) chloride has a very small Ksp, causing it to precipitate readily when chloride ions are added. Copper(II) chloride is highly soluble in water with a large Ksp, so copper ions remain in solution. This difference in solubility enables selective precipitation to separate these metal ions.

A saturated solution contains the maximum amount of dissolved solute at equilibrium, where Q equals Ksp. At this point, solid precipitate exists in dynamic equilibrium with dissolved ions. The solution cannot dissolve more solute without changing conditions like temperature or adding a common ion.

Although AgBr has a smaller Ksp than AgCl, the chloride ion concentration in the mixture was significantly greater than the bromide concentration. AgCl begins to precipitate at a lower silver ion concentration than AgBr because the higher chloride concentration causes Q to exceed Ksp first for the chloride salt.

An unsaturated solution contains dissolved ions where Q is less than Ksp, meaning the solution can dissolve more solute. No solid precipitate forms in an unsaturated solution because ion concentrations remain below their equilibrium values, allowing all ions to remain dissolved.