Solubilty Product Constant: Videos & Practice Problems

What does the solubility product constant (K_sp) represent in the context of an ionic compound?

Given the molar solubility of calcium fluoride (CaF₂) is 2.14 x 10^-4 M, calculate its K_sp.

How does the solubility product constant (K_sp) differ from the general equilibrium constant (K)?

If the reaction quotient (Q) for a dissolution reaction is greater than the solubility product constant (K_sp), what can be inferred about the system?

For a solution of lead(II) iodide (PbI₂), if Q = 1.0 x 10^-8 and K_sp = 7.1 x 10^-9, what will happen?

If the concentration of ions in a solution is increased, how does this affect the reaction quotient (Q) and the equilibrium state?

In a solution where Q = 2.5 x 10^-5 and K_sp = 1.8 x 10^-5 for barium sulfate (BaSO₄), what is the expected outcome?

How does an increase in temperature generally affect the K_sp of an endothermic dissolution reaction?

If a solution of magnesium hydroxide (Mg(OH)₂) is at equilibrium and more Mg²⁺ ions are added, what will happen to the equilibrium position?

Combine your understanding of solubility and K_sp to explain why calcium carbonate (CaCO₃) is less soluble in seawater compared to freshwater.

Why is K_sp considered a special type of equilibrium constant?

In a solution where Q = 3.0 x 10^-6 and K_sp = 2.0 x 10^-6 for a salt, what does this indicate about the system's equilibrium?

For a solution of calcium phosphate (Ca₃(PO₄)₂), if Q = 1.0 x 10^-25 and K_sp = 2.0 x 10^-29, what is the expected outcome?

What condition is described when Q < K_sp for a solution?

If a solution of iron(III) hydroxide (Fe(OH)₃) is at equilibrium and more OH^- ions are added, what will happen to the equilibrium position?

Combine your understanding of solubility and K_sp to explain why gypsum (CaSO₄·2H₂O) is more soluble in acidic solutions.

Why is K_sp used specifically for ionic solids?