Chapter 14, Problem 14.18QP

Chemical Equilibrium II

Magnesium hydroxide. Mg(OH)₂, is a white, partially soluble solid that is used in many antacids. The chemical equation for the dissolving of Mg(OH)₂(s) in water is

${\text{Mg(OH)}}_{\text{2}}(s)\stackrel{}{\rightleftharpoons }{\text{Mg}}^{\text{2+}}(aq)+2{\text{OH}}^{-}(aq)$

1. a Describe a simple experimental procedure that you could use to study this solubility equilibrium. In your experiment, how would you determine when the solution process has attained equilibrium?
2. b Write the equilibrium-constant expression for this dissolving of magnesium hydroxide.
3. c Suppose equilibrium has been established in a container of magnesium hydroxide in water, and you decide to add more solid Mg(OH)₂. What would you expect to observe? What effect will this addition of Mg(OH)₂ have on the concentrations of Mg²⁺(aq) and OH⁻(aq)?
4. d Say you haw prepared an equilibrium solution of Mg(OH)₂ by adding pure solid Mg(OH)₂ to water. If you know the concentration of OH⁻(aq), can you determine the concentration of Mg²⁺(aq)? If not, what information do you need that will allow you to determine the answer?
5. e You slowly add OH⁻ from another source (say, NaOH) to an equilibrium mixture of Mg(OH)₂ and water. How do you expect the concentration of the Mg²⁺(aq) to change? What might you be able to observe happening to the Mg(OH)₂(s) as you add the OH⁻?
6. f Next you remove some, but not all, of the Mg(OH)₂(s) from the mixture. How will this affect the concentrations of the Mg²⁺(aq) and OH⁻(aq)?
7. g If someone hands you a container of Mg(OH)₂(aq) and there is no solid Mg(OH)₂ present, is this solution at equilibrium? If it is not at equilibrium, what could you add to or remove from the container that would give an equilibrium system?
8. h Consider an individual OH⁻(aq) ion in an Mg(OH)₂ solution at equilibrium. If you could follow this ion over a long period of time, would you expect it always to remain as an OH⁻(aq) ion, or could it change in some way?