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Name: Nissana Khan Class: L6 Module 2- Kinetics and Equilibria

Name: Nissana Khan Class: L6 Module 2- Kinetics and Equilibria. Solubility Product. What is solubility product?. The solubility product or K sp is the equilibrium constant for a chemical reaction in which a solid ionic compound dissolves to yield its ions in solution

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Name: Nissana Khan Class: L6 Module 2- Kinetics and Equilibria

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  1. Name: Nissana KhanClass: L6Module 2- Kinetics and Equilibria Solubility Product

  2. What is solubility product? • The solubility product or Ksp is the equilibrium constant for a chemical reaction in which a solid ionic compound dissolves to yield its ions in solution • Many salts which we refer to as insoluble do actually dissolve to a small extent. The term sparingly soluble is usually used for these types of salts. • In a saturated solution of a sparingly soluble salt, equilibrium can be established between dissolved ions and undissolved salt at a given temperature. • For example: BaSO₄ (s) Ba (2+) + SO₄ (2-) (aq) • We can write the equilibrium constant as: Ksp = [Ba(2+) (aq) [SO₄ (2-) (aq)] [BaSO₄ (s)]

  3. What is solubility product? • A constant Ksp multiplied by a constant (the concentration of BaSO₄ (s) ) gives a new constant which we call the solubility product. • Thus the solubility product of BaSO₄ (s) at twenty five degrees celcius is the product of the concentration of the cation and anion in a saturated solution. • Thus the solubility product of a sparingly soluble solution can be defined as the product of the concentrations of the ions in a saturated solution of the salt, raised to the appropriate powers.

  4. Solubility and the Solubility Product Solubility product can be calculated from the solubility of a salt. Example 1: If the solubility of AgCl (s) at 18 ͦC is 1.00 x 10⁻³ mol dm ⁻³, calculate a value for Ksp at this temperature. The solubility means that 1 x 10⁻³ mole of AgCl is dissolved in 1 dm³ of water at equilibrium. From the equation AgCl (s) Ag (+) (aq) + Cl (-) (aq) 1 x 10⁻³ mol dm³ AgCl (s) would form 1 x 10⁻³ mol dm⁻³ Ag (+) ions and 1 x 10⁻³ mol dm⁻³ Cl(-) ions. Ksp = [Ag (+)] [Cl (-) ] = [1 x 10⁻³ mol dm⁻³] [1 x 10⁻³ mol dm⁻³] = 1x 10⁻⁶ mol dm⁻⁶

  5. Solubility and the Solubility Product Limitations of Ksp • Ksp is valid only for saturated solutions in which the concentration of the ions is no more than 0.01 mol dm⁻³. • Ksp is affected by temperature since it is an equilibrium constant.

  6. The common ion effect The common ion effect is an effect which results when two substances, which both ionize to give the same (common) ion, are involved in a chemical equilibrium. • The solubility of a sparingly soluble salt is reduced in a solution that contains an ion in common with that salt. For instance, the solubility of silver chloride in water is reduced if a solution of sodium chloride is added to a suspension of silver chloride in water. • When a reaction has reached equilibrium, and an outside source adds more of one of the ions that is already in solution, the result is to cause the reverse reaction to occur at a faster rate and reestablish the equilibrium. This is called the common ion effect.

  7. The common ion effect BaSO₄ (s) Ba (2+) (aq) + SO₄ (2-) (aq) Na₂SO₄ (s) 2Na (+) (aq) + SO₄ (2-) (aq) The presence of the sodium sulphate increases the concentration of SO₄ ²⁻ ions in the mixture. An increase in the concentration of SO₄ (2-) ions would cause a decrease in concentration of Ba (2+) ions to keep the solubility product as a constant value.

  8. The common ion effect Thus the equilibrium shifts and reduces the solubility of BaSO₄. The addition of the sulphate ions results in the precipitation of solid barium sulphate. The precipitation of a solute on addition of another solution which has an ion in common with the solute is referred to as the COMMON ION EFFECT.

  9. Predicting Precipitation Precipitation of a sparingly soluble salt on addition of another solution which has an ion in common with the salt occurs only if the ionic product of the solution at that moment exceeds the solubility product. To predict whether precipitation of a salt will occur when a solution with a common ion is added we can compare the ionic product of the solution upon mixing to the solubility product. • If the ionic product < Ksp , then no precipitation occurs • If the ionic product = Ksp , then a saturated solution forms • If the ionic product > Ksp , then precipitation occurs

  10. Predicting Precipitation Example 1: Determine if a precipitate will form when a solution of calcium sulphate containing 5 x 10⁻⁴ mol dm⁻³ of Ca (2+) ions is mixed with an equal volume of a solution of 5 x 10⁻⁴ mol dm⁻³ SO₄ (2-) ions at 25 ͦC. The solubility product of CaSO₄ is 2.4 x 10⁻⁵ mol² dm⁻⁶ at 25 ͦC. CaSO₄ (s) Ca (2+) (aq) + SO₄ (2-) (aq) Ksp = [ Ca (2+) (aq) ] [ SO₄ (2-) (aq) ] = 2.4 x 10⁻⁵ mol² dm⁻⁶ Upon mixing and before any precipitation occurs, the [Ca (2+) (aq)] is halved since 5 x 10⁻⁴ mole of Ca (2+) is now present in 2 dm⁻³ of solution.

  11. Predicting Precipitation The [ SO₄ (2-) (aq)] remains the same since twice the number of moles is in 2 dm⁻³ of solution. [ Ca (2+) (aq) ] = 2.5 x 10⁻⁴ mol dm⁻³ [ SO₄ (2-) (aq) ] = 5 x 10⁻⁴ mol dm⁻³ The ionic product upon mixing = [Ca (2+)] [ SO₄ (2-)] = ( 2.5 x 10⁻⁴) ( 5 x 10⁻⁴) = 1.25 x 10⁻⁷ mol² dm⁻⁶ Since Ksp = 2.4 x 10⁻⁵ mol² dm⁻⁶ , the ionic product upon mixing is less than the solubility product therefore no precipitation occurs.

  12. References • Chemistry for CAPE by Susan Maraj and Arnold Samai • Wikipedia, the free encyclopedia • Wisegeek.com • ChemGuide.com

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