Solubility product, commonly written as Ksp. It describes the equilibrium that exists when an ionic solid dissolves slightly in water. Even substances considered "insoluble" dissolve to a tiny extent, and Ksp quantifies exactly how much. Understanding Ksp helps you predict whether a precipitate will form, calculate solubility, and compare the stability of ionic compounds in solution.
What Is Solubility Product (Ksp)?
Ksp (solubility product constant) is the equilibrium constant for the dissolution of a sparingly soluble ionic solid.
It represents:
How much of the solid dissolves
The concentration of ions at equilibrium
The point at which the solution becomes saturated
Even for nearly insoluble compounds, Ksp has a definite value.
The Dissolution Equilibrium
When a sparingly soluble salt dissolves, an equilibrium forms between the solid and its ions.
Example: AgCl(s) ⇌ Ag⁺(aq) + Cl⁻(aq)
At equilibrium, the concentrations of Ag⁺ and Cl⁻ remain constant, and the solid no longer dissolves or forms crystals overall.
The Ksp expression is:
Ksp = [Ag⁺][Cl⁻]
The solid does not appear in the expression because its concentration is constant.
Learn why ionic and covalent bonds form under different conditions and how electronegativity, energy and structure determine bonding type.
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3. Fe(OH)₃
Fe(OH)₃(s) ⇌ Fe³⁺ + 3OH⁻ Ksp = [Fe³⁺][OH⁻]³
Larger exponents make the Ksp expression highly sensitive to ion concentrations.
What Ksp Tells You
Ksp shows how soluble a substance is:
Larger Ksp → more soluble
Smaller Ksp → less soluble
However, Ksp values are temperature-dependent. Increasing temperature usually increases solubility of solids.
Calculating Molar Solubility from Ksp
Molar solubility (s)
The number of moles of solute that dissolve per liter of solution.
For 1:1 salts (AgCl, NaF):
Ksp = s²
s = √Ksp
For 1:2 or 1:3 salts, use stoichiometry:
Example: CaF₂ [s] = concentration of Ca²⁺ [2s] = concentration of F⁻ Ksp = (s)(2s)² = 4s³ Solve for s.
IB exam questions frequently require these calculations.
Predicting Precipitation Using Ksp
Ksp helps determine whether a precipitate will form when two solutions are mixed.
You use Q, the reaction quotient:
If Q < Ksp → no precipitate
If Q = Ksp → saturated solution
If Q > Ksp → precipitate forms
This is vital in qualitative analysis and selective precipitation techniques.
Example: Mixing AgNO₃ and NaCl forms a precipitate when: [Ag⁺][Cl⁻] > Ksp of AgCl
This ability to predict precipitation appears often on IB Paper 2.
Common Uses of Ksp in IB Chemistry
1. Determining solubility
Use Ksp to calculate how much of a salt dissolves.
2. Predicting precipitate formation
Compare Q to Ksp.
3. Removing ions from solution
Industries use precipitation to remove impurities.
4. Selective precipitation
Different Ksp values allow chemists to separate ions selectively.
5. Acid–base influence on solubility
Hydroxides and carbonates become more soluble in acidic conditions because acid removes an ion involved in equilibrium.
Factors Affecting Solubility and Ksp
1. Temperature
Ksp changes with temperature; most solids become more soluble when heated.
2. Common-ion effect
Adding an ion already in the equilibrium reduces solubility.
Example: Adding NaCl reduces AgCl solubility because Cl⁻ increases.
3. pH
Hydroxide and carbonate salts dissolve more readily in acidic solutions.
Common IB Misunderstandings
“A small Ksp means no ions dissolve.”
Incorrect—small Ksp means few ions dissolve, not zero.
“Ksp and solubility are the same.”
Ksp is a constant; solubility is a numerical amount.
“A large Ksp always means high solubility.”
Generally true, but stoichiometry matters.
“Solids appear in equilibrium expressions.”
Solids never appear in Ksp expressions.
FAQs
Why doesn’t the solid appear in the Ksp expression?
Its concentration is constant and does not change during equilibrium.
Is Ksp affected by concentration changes?
No—only temperature affects Ksp. Concentration affects Q.
Why do some salts dissolve more in acid?
Acid removes one of the ions (like OH⁻ or CO₃²⁻), shifting equilibrium.
Conclusion
The solubility product constant (Ksp) describes the equilibrium between a sparingly soluble ionic solid and its dissolved ions. It allows IB Chemistry students to quantify solubility, predict precipitation, and understand the factors that influence solution equilibria. Mastering Ksp is essential for solving advanced equilibrium questions and performing accurate stoichiometric calculations.
