mc_ch13

Objectives for Chapter 13

  • Write equations for the dissolution of soluble ionic compounds in water.

  • Predict whether a precipitate will form when solutions of soluble ionic compounds are combined, and write net ionic equations for precipitation reactions.

  • Compare the dissociation of ionic compounds with the ionization of molecular compounds.

Objectives Continued

  • Draw the structure of the hydronium ion, and explain why it is used to represent the hydrogen ion in solution.

  • Distinguish between strong electrolytes and weak electrolytes.

Dissociation

  • Dissociation is the separation of ions that occurs when an ionic compound dissolves in water.

  • Example: NaCl dissociates into Na⁺ and Cl⁻.

Dissociation of NaCl

  • Sodium chloride (NaCl) dissociates into Na⁺ and Cl⁻ in a 1:1 ratio when dissolved in water.

  • The volume of the solution does not affect the dissociation ratio.

Sample Problem A: Aluminum Sulfate

  • Write the dissolution equation for aluminum sulfate, Al₂(SO₄)₃.

    • Dissolution: Al₂(SO₄)₃ → 2Al³⁺ + 3SO₄²⁻

  • Moles produced upon dissolving 1 mol of aluminum sulfate: 2 moles of Al³⁺ and 3 moles of SO₄²⁻, totaling 5 moles of ions.

Precipitation Reactions

  • Compounds of very low solubility can be considered insoluble for practical applications.

General Solubility Guidelines

  • Sodium, potassium, and ammonium compounds are soluble in water.

  • Nitrates, acetates, and chlorates are soluble.

  • Most chlorides are soluble except for those of silver, mercury(I), and lead.

  • Most sulfates are soluble except for those of barium, calcium, mercury, strontium, and lead.

  • Most carbonates, phosphates, and silicates are insoluble except for sodium, potassium, and ammonium.

  • Most sulfides are insoluble except for calcium, strontium, sodium, potassium, and ammonium.

Soluble and Insoluble Ionic Compounds

  • Soluble in water: NiCl₂, KMnO₄, CuSO₄, Pb(NO₃)₂

  • Insoluble in water: AgCl, CdS

Formation of a Precipitate

  • Example: NaCl dissociates in water, AgNO₃ does the same. When combined, the reaction:

  • Ag⁺(aq) + Cl⁻(aq) → AgCl(s) precipitate.

Net Ionic Equations

  • Net ionic equations include only those compounds and ions that undergo a chemical change.

  • Spectator ions do not participate in the reaction.

Writing a Net Ionic Equation

  • Example: K₂SO₄(aq) + Ba(NO₃)₂(aq) → 2KNO₃(aq) + BaSO₄(s)

  • Total ionic equation: K⁺(aq) + SO₄²⁻(aq) + Ba²⁺(aq) + 2NO₃⁻(aq) → 2K⁺(aq) + 2NO₃⁻(aq) + BaSO₄(s)

  • Net ionic equation: SO₄²⁻(aq) + Ba²⁺(aq) → BaSO₄(s)

Ionization Explained

  • Ionization is the formation of ions from solute molecules through solvent action.

  • Example: HCl ionizes to H⁺ and Cl⁻ in aqueous solution.

  • The hydronium ion (H₃O⁺) represents the hydrogen ion in solution.

Strong and Weak Electrolytes

  • Electrolytes: substances yielding ions to conduct an electric current.

  • Strong electrolytes: fully dissociate into ions (e.g., HCl, NaCl).

  • Weak electrolytes: partially dissociate (e.g., HF).

Colligative Properties of Solutions

  • Colligative properties depend on the concentration of solute particles, not their identity.

  • Key colligative properties include:

  • Vapor-Pressure Lowering

  • Freezing-Point Depression

  • Boiling-Point Elevation

  • Osmotic Pressure

Vapor-Pressure Lowering

  • Nonvolatile substances lower the vapor pressure compared to the pure solvent.

Freezing-Point Depression

  • Freezing-point depression ( ∆tf) is directly proportional to the molal concentration of the solution, where ∆tf = Kf * m.

Boiling-Point Elevation

  • Boiling-point elevation ( ∆tb) is also directly proportional to molal concentration: ∆tb = Kb * m.

Osmotic Pressure

  • Osmosis occurs through a semipermeable membrane, moving solvent from lower to higher solute concentration.

  • Osmotic pressure is the pressure required to stop osmosis.