Gen Chem 1 ~ 1112

Chapter 4: Solubility Rules

Soluble Ionic Compounds

  • Compounds containing:

    • NO₃⁻: none

    • CH₃COO⁻: none

    • Cl⁻: exceptions include Ag⁺, Hg₂²⁺, Pb²⁺

    • Br⁻: exceptions include Ag⁺, Hg₂²⁺, Pb²⁺

    • I⁻: exceptions include Ag⁺, Hg₂²⁺, Pb²⁺

    • SO₄²⁻: exceptions include Sr²⁺, Ba²⁺, Hg₂²⁺, Pb²⁺

Insoluble Ionic Compounds

  • Compounds containing:

    • S²⁻: exceptions include NH₄⁺, alkaline metal cations (Ca²⁺, Sr²⁺, Ba²⁺)

    • CO₃²⁻: exceptions include NH₄⁺, alkaline metal cations

    • PO₄³⁻: exceptions include NH₄⁺, alkaline metal cations

    • OH⁻: exceptions include NH₄⁺, alkaline metal cations (Ca²⁺, Sr²⁺, Ba²⁺)

Exchange Reaction

  • General form: AB + CD → AD + CB

Precipitation Reactions

  • Occurs if one or both products (AD + CB) are insoluble in water.

  • If both products are soluble in water, no precipitation occurs (no reaction).

Steps to Write a Net Ionic Equation:

  1. Write the full and balanced molecular equation.

  2. Use solubility rules to determine which products are soluble and which are insoluble (precipitates).

  3. Create a complete ionic equation, placing (aq) next to soluble ions and (s) next to all precipitates.

  4. Cancel out spectator ions (ions that appear on both sides of the equation).

  5. Rewrite the complete ionic equation without spectator ions; this is the net ionic equation.

Strong Acids and Bases

Strong Acids

  • Acids that completely dissociate in water to form H⁺:

    • Hydrochloric Acid (HCl)

    • Hydrobromic Acid (HBr)

    • Hydroiodic Acid (HI)

    • Perchloric Acid (HClO₄)

    • Nitric Acid (HNO₃)

    • Sulfuric Acid (H₂SO₄)

Strong Bases

  • Bases that completely dissociate in water to form OH⁻:

    • Group IA Metals - OH: LiOH, NaOH, KOH, RbOH, CsOH

    • Heavy Group IIA Metals - (OH)₂: Ca(OH)₂, Sr(OH)₂, Ba(OH)₂

Acid-Base (Neutralization) Reactions

  • Neutralization reactions between acids and bases always produce water (H₂O) and salt.

  • General form: AB + CD → AD + CB

Oxidation-Reduction (Redox) Reactions

Definitions

  • Oxidation: Losing electrons.

  • Reduction: Gaining electrons.

Oxidation Number Rules:

  1. An atom in elemental form has an oxidation number of zero (e.g., H₂, Na).

  2. For monatomic ions, the oxidation number equals the ion's charge (e.g., Mg²⁺ = +2).

  3. Nonmetals usually have negative oxidation numbers but can be positive.

    • Oxygen: -2 (except in peroxides).

    • Hydrogen: +1 when bonded to nonmetals, -1 when bonded to metals.

    • Fluorine: always -1.

    • Halogens are usually -1 but can be positive if bonded to oxygen.

Steps to Determine Oxidation in a Redox Reaction:

  1. Determine the oxidation number of every element in the chemical equation.

  2. Elements whose oxidation number increases (becomes more positive) are oxidized (reducing agents).

  3. Elements whose oxidation number decreases (becomes more negative) are reduced (oxidizing agents).

Concentrations in Solutions

Molarity (M)

  • A way of expressing a solution's concentration:

    • Formula: M = Moles of solute / Volume of solution in liters

    • Units: moles/liter

Steps for Molarity Problems:

  1. Remember: Molarity units are moles/liter.

  2. Look at what is given, especially units.

  3. Determine what is being asked and what units the final answer should be in.

  4. Use dimensional analysis, focusing on the units.

Example Problem:

  • How many moles of HCl are present in 35.0 mL of a 4.50 M solution?

    • Calculation:

      • (35.0 mL) * (4.50 mol / 1000 mL) = 0.158 mol HCl.