Untitled Quiz

Solubility Equilibria

• Solubility equilibria describes the dynamic equilibrium between a solid solute and its dissolved ions in a solution.
• A saturated solution exists when the rates of dissolution and crystallization are equal.
• The concentration of a saturated solution is its solubility.
• Solubility varies with temperature.
• A solubility-temperature graph is a solubility curve.
• Most ionic substances' solubility increases with temperature.
• Exceptions include anions such as SO₃^2-, SO₄^2-, AsO₄^3-, and PO₄^3-.
• An unsaturated solution contains less solute than a saturated solution.
• A supersaturated solution contains more solute than a saturated solution. Supersaturated solutions are unstable and will precipitate.

Saturated Solutions

• The concentration of a saturated solution is called its solubility.
• Solubility varies with temperature.
• A solubility-temperature graph is called a solubility curve.
• Solubility for various salts is shown on graphs, each line representing a different substance.

Aqueous Solubility

• Aqueous solubility is the amount of solute that can dissolve in 100 g of water at a given temperature.
• Solubility can be expressed in various ways such as molarities, mass percentages, or grams of solute/100 g of water.
• Points on the solubility curve represent saturated solutions. Areas above the curve represent supersaturated solutions and below unsaturated solutions.

Solubility

• Generally, solubilities of ionic substances increase with increasing temperature.
• Compounds containing anions (SO₃^2-, SO₄^2-, AsO₄^3-, and PO₄^3-) are exceptions.
• Adding solute to a solvent, with less than the amount needed for saturation, can result in an unsaturated solution, which the additional solute will dissolve.
• A saturated solution can be prepared by dissolving the solute at a high temperature and then allowing it to cool slowly.

Supersaturated Solutions

• Usually, excess solute crystallizes from a solution, but occasionally all the excess solute may stay in solution.
• Supersaturated solutions are unstable.
• If crystals of solute are added, the excess solute crystallizes.

Le Châtelier's Principle

• A system at equilibrium subjected to a change will spontaneously shift to counteract the change.
• Adding a product causes the equilibrium to shift to the left (reactants).
• Adding a reactant causes the equilibrium to shift to the right (products).

The Solubility Product Constant, K_sp

• The equilibrium constant for the reaction in which a solid salt dissolves to yield its constituent ions in solution.
• The concentration of the solid is omitted from the equilibrium constant expression because it is in its standard state.
• Example: CaSO₄ (s) ⇌ Ca²⁺(aq) + SO₄^2-(aq).
• K_sp = [Ca²⁺] [SO₄^2-]
• K_sp = 9.1 × 10⁻⁶ (at 25°C)

Solubility vs. Solubility Product

• Solubility is an equilibrium position for a solution; a specific solubility value for a salt at particular temperatures.
• The solubility product, K_sp, is a constant value for a specific salt, at any specific temperature. The solubility product depends on ion activities.

The Relationship Between Solubility and K_sp

• There's a relationship between the solubility product constant (K_sp) and the molar solubility of a solute.
• Molar solubility is a solute's molarity in a saturated aqueous solution.

Relative Solubilities

• Compare K_sp values only for salts that have the same number of ions (e.g., AgCl, Cul).
• The larger the K_sp value, the more soluble the salt.

Fractional Precipitation

• Fractional precipitation is a technique used to separate two or more ions that can be precipitated by the same reagent.
• The rate of reagent addition is controlled to selectively precipitate different ions in a solution.
• The slower the reagent addition, the greater the chance the less soluble of the two products will precipitate first.

Solubility and pH

• The pH of a solution can greatly affect the solubility of a salt.
• This is especially true for salts whose anions are the conjugate bases of weak acids or the base OH⁻ itself.
• An example is Mg(OH)₂, which is highly insoluble in water and is a common antacid commonly known as milk of magnesia.
• Hydroxide ions from magnesium hydroxide react with hydronium ions from stomach acid to form water.
• Increasing acidity increases the solubility of salts containing basic anions.

Incomplete Dissociation of Solute into Ions

• In solubility calculations, assume complete dissociation into ions, but this is often not a valid assumption.
• Solute might not be 100% ionic and some may dissolve as molecules or as ion pairs.

Simultaneous Equilibria

• The reversible reaction between a solid solute and its ions in solution, is never the only process.
• Other simultaneous processes occurring include water autoionization and reactions occurring between solute ions and other solution species. These may include acid-base reactions and complex ion formation.

Limitations of K_sp

• K_sp values are often limited to slightly soluble or essentially insoluble solutes.
• When calculating K_sp it is important to consider if other equilibria are occurring simultaneously
• Experimental solubility values can differ from calculated K_sp values, as the calculated ones are based on assuming complete dissociation.
• Molarities are often used in place of activities to simplify calculations.