On the Solubility Product of Lead Acetate
I have heard that all substances have their own solubility in water. In the case of lead acetate, there is a state of dissolution equilibrium in the solution.

The dissolution equilibrium is when the dissolution rate of the solute is equal to the crystallization rate. When lead acetate is placed in water, it partially dissociates into lead ions and acetate ions, but its dissociation is not endless.

The solubility product of lead acetate ($K_ {sp} $) is of great significance. $K_ {sp} $is a constant product of the concentration powers of each ion in a saturated electrolyte solution at a certain temperature. For lead acetate ($Pb (CH_ {3} COO) _ {2} $), the solution equilibrium expression is $Pb (CH_ {3} COO) _ {2} (s) \ rightleftharpoons Pb ^ {2 +} (aq) + 2CH_ {3} COO ^ {-} (aq) $, then its solubility product $K_ {sp} = c (Pb ^ {2 +}) \ times c ^ {2} (CH_ {3} COO ^{-})$。

If the $K_ {sp} $of lead acetate at a certain temperature is known, it can be judged whether the solution is saturated according to the concentration of lead ion and acetate ion in the solution. If $Q_ {c} $ (ion product, calculated in the same way as $K_ {sp} $, but the concentration calculation value at any time) is less than $K_ {sp} $, the solution is not saturated, and lead acetate can continue to dissolve; if $Q_ {c} $is equal to $K_ {sp} $, the solution is saturated, and the dissolution is equal to the crystallization rate; if $Q_ {c} $is greater than $K_ {sp} $, lead acetate will crystallize.

To measure the $K_ {sp} $of lead acetate, the solution containing lead ions and acetate ions can be carefully configured, and the concentration of the two ions at equilibrium can be accurately determined by clever experimental means, and then $K_ {sp} $can be calculated according to the above formula. This process requires rigorous operation, so that the error is not too large, in order to obtain the accurate $K_ {sp} $value, so that the lead acetate dissolves at this temperature.