Ch 4 - Stoichiometry of Chemical Reactions

Chapter 4: Stoichiometry of Chemical Reactions

Chapter Outline

  • 4.1 Writing and Balancing Chemical Equations

  • 4.2 Classifying Chemical Reactions

  • 4.3 Reaction Stoichiometry

  • 4.4 Reaction Yields

  • 4.5 Quantitative Chemical Analysis

4.1 Writing and Balancing Chemical Equations

  • Chemical Equation Basics:

    • Previous chapters introduced element symbols representing atoms, molecules, and compounds.

    • A balanced chemical equation symbolizes identities and relative quantities of substances in chemical/physical changes.

    • Example: Modern rocket fuels are mixtures measured for thrust-generating reactions.

  • Chemical Representation of Reactions:

    • Example: Methane (CH₄) and oxygen (O₂) react to form carbon dioxide (CO₂) and water (H₂O).

  • Writing Chemical Equations:

    1. Reactants: Substances undergoing a reaction, placed on the left side.

    2. Products: Generated substances, placed on the right side.

    3. Symbols: Plus signs (+) separate reactants/products; an arrow (⟶) separates reactant and product sides.

    4. Coefficients: Represent relative amounts; a coefficient of 1 is typically omitted.

  • Coefficients Explanation:

    • Use the smallest whole-number coefficients, indicating relative numbers of reactants/products as ratios.

    • Example ratio for methane and oxygen reacting is 1:2:1:2.

  • Balancing Chemical Equations:

    • A balanced equation shows equal numbers of each atom on reactant and product sides.

    • Atom counts are determined by multiplying coefficients by subscripts in formulas.

  • Balancing Process Steps:

    • Compute atoms for elements appearing in multiple formulas.

    • Example: Oxygen appears in both CO₂ and H₂O in the methane reaction.

  • Example Balancing for Methane Reaction:

    • The balanced reaction: CH₄ + 2O₂ ⟶ CO₂ + 2H₂O.

    • Verification of balance:

    • Carbon: 1 C from reactants and products.

    • Hydrogen: 4 H from reactants; 4 H in products.

    • Oxygen: 4 O from reactants; 4 O from products.

  • Balancing by Inspection:

    • Example: Unbalanced H₂O ⟶ H₂ + O₂.

    • Check balance for H and O; adjust coefficients to achieve balance.

  • Steps for Balancing Chemical Equations:

    1. Write correct formulas for reactants (left) and products (right).

    • Example: Ethane (C₂H₆) reacts with O₂ to yield CO₂ and H₂O.

    1. Adjust coefficients to balance atoms without changing subscripts—C₂H₆ + O₂ ⟶ CO₂ + H₂O.

    2. Start balancing elements appearing in one reactant/product first.

    • Example: Balance C, then H, followed by O.

    1. Balance elements appearing in multiple species.

    2. Check totals for atoms to ensure balance.

  • Using Fractions in Balancing:

    • Fractional coefficients can simplify balancing; multiply by whole numbers to eliminate fractions.

    • Example: From C₂H₆ + O₂ ⟶ 3H₂O + 2CO₂ to 2C₂H₆ + 7O₂ ⟶ 4CO₂ + 6H₂O.

  • Additional Information in Equations:

    • Physical states often indicated with abbreviations:

    • (g) for gas, (l) for liquid, (s) for solid, (aq) for aqueous.

    • Special conditions (e.g., heating) may be noted above/below the equation’s arrow (e.g., reaction at Δ).

    • Example: CaCO₃(s) ⟶ CaO(s) + CO₂(g).

  • Ionic Reactions:

    • Many reactions occur in aqueous media, expressed at various detail levels, including:

    • Molecular equation: CaCl₂(aq) + 2AgNO₃(aq) ⟶ Ca(NO₃)₂(aq) + 2AgCl(s).

    • Complete ionic equation: Representing all dissociated ions:

    • Example: Dissociation: CaCl₂(aq) ⟶ Ca²⁺(aq) + 2Cl⁻(aq).

  • Net Ionic Equations:

    • Spectator ions do not change in a reaction.

    • Eliminate spectator ions—resulting in a net ionic equation:

    • Ag⁺(aq) + Cl⁻(aq) ⟶ AgCl(s).

4.2 Classifying Chemical Reactions

  • Precipitation Reaction:

    • Dissolved substances react to form one or more solid products (also double displacement reactions).

    • Involves ion exchange in ionic compounds in aqueous solutions.

  • Solubility and Precipitation:

    • Solubility: Maximum concentration achievable under specified conditions.

    • Soluble substances: high solubility.

    • Insoluble substances: low solubility that readily precipitate from solutions.

  • Common Ionic Compounds Solubility:

    • Soluble Compounds:

    • Group 1 cations and ammonium ions.

    • Halides, acetates, bicarbonates, nitrates, chlorates, and sulfates (with exceptions for certain metals).

    • Insoluble Compounds:

    • Carbonates, chromates, phosphates, sulfides, hydroxides (exceptions for group 1 cations and Ba²⁺).

  • Precipitation Reaction Examples:

    • Molecular Equation:

    • 2KI(aq) + Pb(NO₃)₂(aq) ⟶ PbI₂(s) + 2KNO₃(aq).

    • Ionic Equation: