Aqueous Equilibrium Notes

Solubility of Ionic Compounds

• General Properties

  • Ionic compounds often act as strong electrolytes when dissolved in water.

  • Definition: Strong electrolytes fully dissociate into ions in water.

  • Example:

  • [ \text{NaCl} \rightarrow \text{Na}^+ + \text{Cl}^- ]

  • Not all ionic compounds have infinite solubility; they reach a limit at saturation.

• Saturated Solution:

  • The maximum amount of solute that can dissolve in a solution under specified conditions.

  • Many salts completely dissolve, while others may be poorly soluble, classifying them as insoluble or sparingly soluble.

Solubility Rules for Ionic Compounds

• General Solubility:

  • Compounds with the following ions are typically soluble:

  • Li⁺, Na⁺, K⁺, NH₄⁺

  • NO₃^-, C₂H₃O₂^-

  • Cl^-, Br^-, I^- (with exceptions for Ag⁺, Hg₂²⁺, Pb²⁺)

  • SO₄^2- except with Ba²⁺, Sr²⁺, Pb²⁺

• General Insolubility:

  • Compounds containing OH^-, S^2-, CO₃^2-, PO₄^3- are usually insoluble, unless paired with soluble ions like Li⁺, Na⁺, K⁺, NH₄⁺.

Specific Examples in Solubility

• Lead(II) Chloride (PbCl₂):

  • Characteristics:

  • Considered a strong electrolyte, yet sparingly soluble in water.

  • Reaction:

    • Dissolves as follows:

    • [ \text{PbCl}_2(s) \rightleftharpoons \text{Pb}^{2+}(aq) + 2\text{Cl}^-(aq) ]

  • Combined equilibrium representation:

    • [ \text{PbCl}_2(s) \rightleftharpoons \text{Pb}^{2+}(aq) + 2\text{Cl}^-(aq) ]

• General Equilibrium for Ionic Compounds:

  • For a compound ( ext{MnXm} ), the equilibrium can be expressed as:

  • [ \text{MnXm}(s) \rightleftharpoons n \text{Mm}^+(aq) + m \text{Xn}^-(aq) ]

  • Calculate K_sp as follows:

  • [ K{sp} = [\text{Mm}^+]^n{eq} [\text{Xn}^-]^m_{eq} ]

Common Ion Effect

• Understand that introducing a common ion to a solution reduces the solubility of an ionic compound.

• Derived from Le Chatelier's principle; increases equilibrium concentration of the products, thereby shifting the dissolution equilibrium towards the solid phase, decreasing its solubility.

pH Effect on Solubility

• For Ionic Hydroxides (e.g., Ca(OH)₂):

  • Equalities and effects of pH:

  • [ ext{Ca(OH)}_2(s) \rightleftharpoons ext{Ca}^{2+}(aq) + 2 \text{OH}^-(aq) ]

  • Raising pH (more OH^-): decreases solubility

  • Lowering pH (more H⁺): increases solubility

• For Weakly Basic Anions (e.g., CO₃^2-):

  • Example:

  • [ ext{FeCO}3(s) \rightleftharpoons ext{Fe}^{2+}(aq) + ext{CO}3^{2-}(aq) ]

  • A lower pH increases the solubility of the basic ionic compound.

Precipitation of Ionic Compounds

• A salt precipitates when the concentration of cations and anions exceeds that of the solubility limit.

• To Determine Precipitation:

  • Calculate Q_sp and compare with K_sp:

  • If Q_sp = K_sp: solution is saturated / at equilibrium.

  • If Q_sp < K_sp: unsaturated, no precipitation.

  • If Q_sp > K_sp: supersaturated, precipitation occurs.

• Selective Precipitation:

  • Requires that the K_sp of the target insoluble salt is at least 10³ times smaller than other potential ionic compounds in solution.

• Next Class: March 27

• Topics:

  • Chapter 18 – Thermodynamics

  • Discuss Entropy and Enthalpy

  • Homework #7 Due: March 28

  • 10 questions, no time limit, one attempt available before due date; unlimited attempts will be opened after for review.