Topics 7.11 - 7.12

Unit 7: Equilibrium Overview of Solubility Equilibria and the Common Ion Effect

Key Topics for Understanding Solubility Equilibria

• Chemical Equilibrium: This concept refers to a state in which the concentrations of reactants and products remain constant over time due to a reversible reaction. In solubility, it involves the balance between the dissolution of solid salts and the precipitation of ions back into the solid form.

• Le Chatelier's Principle: This principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifts to counteract the change. For solubility, this means that if you increase the concentration of a solute or change the temperature, the dissolution process will adjust. For instance, a decrease in temperature might favor the formation of more solid precipitate.

• Ionic Compounds: Understanding insoluble and soluble ionic compounds is essential for predicting solubility. These compounds dissociate in water to form ions; knowing which ions are produced helps in understanding their interactions and solubility behaviors.

• Molarity: This measurement of concentration is defined as the number of moles of a solute per liter of solution (mol/L). Molarity is used in calculating solubility product constants (Ksp) and is crucial in laboratory settings for preparing solutions.

• Stoichiometry: Fundamental to solving chemical equations, stoichiometry involves using the ratios of the coefficients from balanced equations to relate the amounts of reactants and products, including ions in dissolution reactions. This is critical when calculating Ksp values based on the amounts of dissolved ions.

• Ksp (Solubility Product Constant): This equilibrium constant quantifies the solubility of ionic compounds in water, giving insights into the maximum concentration of ions in a saturated solution at a specific temperature. Generally, a Ksp greater than 1 indicates high solubility, while a Ksp lower than 1 suggests low solubility.

Introduction to Solubility Equilibria

• Ksp (Solubility Product Constant): Ksp gives a numeric representation of the solubility of an ionic compound, indicating the limit of how much can dissolve at a set temperature. Different compounds and temperatures will yield different Ksp values, making it an essential part of solubility equations.

• Solubility Rules and Ksp: These rules help categorize salts: for example, salts with alkali metals (like sodium or potassium) or ammonium ions tend to be soluble, while those with heavy metals often are not. Understanding these rules is key to predicting whether a salt will dissolve.

• Precipitation Reactions: This process occurs when a reaction produces an insoluble compound from dissolved ionic species, leading to the formation of a solid precipitate, which may cloud the solution. Commonly soluble ions include sodium (Na+), potassium (K+), ammonium (NH4+), and nitrate (NO3-), while others may not dissolve readily.

• Dynamic Equilibrium: In a saturated solution, the number of ions entering the solution equals the number of ions leaving it, maintaining equilibrium. The Ksp is derived from the concentrations of the dissolved ions while the solid precipitate is excluded from the expression since its concentration remains constant.

Examples of Dissolution:

• Silver Chloride:

  • Reaction: AgCl(s) ⇌ Ag+(aq) + Cl-(aq)

  • Ksp = [Ag+][Cl-]

• Lead(II) Iodide:

  • Reaction: PbI2(s) ⇌ Pb2+(aq) + 2I-(aq)

  • Ksp = [Pb2+][I-]^2

• Silver Chromate:

  • Reaction: Ag2CrO4(s) ⇌ 2Ag+(aq) + CrO4^2-(aq)

  • Ksp = [Ag+]^2[CrO4^2-]

Common Compound Examples:

• Chromium(III) Hydroxide:

  • Ksp = [Cr3+][OH-]^3

• Magnesium Phosphate:

  • Ksp = [Mg2+]^3[PO4^3-]^2

Important Differences:

• Ksp vs. Solubility: Ksp is dimensionless and indicates the solubility equilibrium constant. Solubility indicates how much of a substance can dissolve, described in terms of molarity (mol/L) or grams per liter.

Common Ion Effect

• The Common Ion Effect illustrates how the addition of an ion common to a dissolved salt decreases its solubility. This happens because the equilibrium shifts to reduce the effect of the added ion, demonstrating Le Chatelier's principle in action.

Example Problem - Calcium Hydroxide Dissolution:

• Reaction: Ca(OH)2(s) ⇌ Ca2+(aq) + 2OH-(aq)

• Given Ksp = 5.0 × 10^-6 shows that introducing more Ca2+ or OH- from another source will reduce the solubility of Ca(OH)2 as the equilibrium shifts to counteract the increase.

Impact of Volume Change:

• When a saturated solution undergoes evaporation or dilution, the concentrations of the ions largely remain unchanged, which is crucial for predicting how solubility might shift under varying conditions.

Comparison of Salts:

• By examining Ksp values, one can compare the relative solubility of different salts without needing to determine their specific ratios in solution, aiding in predicting precipitation reactions.

Precipitate Formation and Qsp:

• To determine precipitate formation, compare the reaction quotient (Qsp) to Ksp:

  • If Qsp < Ksp: No precipitate forms, signifying that more solid can still dissolve.

  • If Qsp > Ksp: A precipitate forms, indicating that solubility limits have been reached.

Common Ion Example:

• In a scenario where sodium sulfate is mixed with barium nitrate, the addition of sulfate ions may result in the precipitation of barium sulfate, a less soluble compound.

Practice Problems for Ksp and Solubility:

• Students should practice by writing dissolution equations and calculating Ksp expressions for various salts. They should work on problems that involve determining Ksp from given solubility values and vice versa, emphasizing real-world applications of solubility and equilibrium concepts.