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General Chemistry I - Chemical Reactions and Stoichiometry

CHY-103 General Chemistry I - Chapter 4: Chemical Reactions and Stoichiometry

Lecture Overview

  • Date: September 11, 2025

  • Location: Toronto Metropolitan University

  • Department: Chemistry & Biology, Faculty of Science

Topics Covered:

  • Balancing Chemical Equations

  • Solutions and Solubility

  • Ionic Solutions

  • Solubility of Ionic Compounds

  • Net Ionic Equations

  • Types of Chemical Reactions

    • Precipitation

Note: Not responsible for balancing redox equations on pages 121-124.

Balancing Chemical Equations

  • Definition: A balanced chemical equation has the correct coefficients for all reactants and products, ensuring mass conservation.

  • Conservation of Mass: Atoms are neither created nor destroyed in a chemical reaction.

General Notes:

  • Orientation: Reactants are located on the left side of the equation, while products are on the right side.

  • Avoiding Fractions: Chemical equations should be presented with the smallest set of whole-number coefficients.

  • States of Matter: States of reactants/products are denoted as:

    • (s) - solid

    • (l) - liquid

    • (g) - gas

    • (aq) - aqueous

  • Adjustment: Only the coefficients are adjustable during balancing; subscripts must remain unchanged.

  • Reaction Conditions: Indicated above the reaction arrow.

Example:

  • Combustion of Propane:
    ext{C}3 ext{H}8(g) + 5 ext{O}2(g) ightarrow 3 ext{CO}2(g) + 4 ext{H}_2 ext{O}(l)

Concept Check 4.1

  • Balance the Reaction:
    ext{Ca}3( ext{PO}4)2 + ext{H}3 ext{PO}4 ightarrow ext{Ca}( ext{H}2 ext{PO}4)2
    Answer Choices:

  • A. 3, 4, 1

  • B. 1, 4, 3

  • C. 1, 3, 4

  • D. 4, 3, 1

Solutions and Solubility

  • Definition of a Solution: A homogeneous mixture of two or more substances, consisting of a solvent and solute.

    • Solvent: The component present in the greatest quantity, often in the same state as the solution (aqueous if water).

    • Solute: The substance(s) that dissolve within the solvent.

Interaction Forces:

  • When a solute dissolves in a solvent, solute-solvent interactions dominate over solute-solute and solvent-solvent interactions.

  • Example interactions:

    • Sodium Chloride (NaCl) Dissolving

Ionic Solutions

  • Electrical Conductivity:

    • Electrolytes: Substances that dissociate into ions in solution and conduct electricity, like most ionic compounds.

    • Examples of Electrolytes:

    • Strong Electrolytes: Almost completely ionize in solution (e.g., strong acids and bases).

    • Weak Electrolytes: Partially ionize in solution (e.g., weak acids).

Examples:

  • Nonelectrolyte:
    ext{C}{12} ext{H}{22} ext{O}_{11}(aq) (Sucrose)

  • Weak Electrolyte:
    ext{CH}_3 ext{COOH}(aq) (Acetic Acid)

  • Strong Electrolyte:
    ext{NaCl}(aq)

Classification of Electrolytes:
Strong Electrolytes

  • Examples:

    • HCl, HBr, HI

    • HNO3, H2SO4, NaOH, KOH

    • Soluble ionic compounds

Weak Electrolytes

  • Examples:

    • HF, CH3OH (Methyl Alcohol), C2H5OH (Ethyl Alcohol)

Nonelectrolytes

  • Examples:

    • Most organic compounds, e.g., Sucrose (C12H22O11)

Note: Do not confuse solubility with electrolyte classification.

Solubility of Ionic Compounds

  • Definition: The ability of a substance to dissolve in a solvent, termed solubility.

    • Soluble: Substances that dissolve readily.

    • Insoluble: Substances with limited or no solubility.

Predicting Solubility:

  • General Solubility Rules for Ionic Compounds in Water:

  1. Salts with Group 1 metals (Li+, Na+, K+) and ammonium ions (NH4+) are soluble.

  2. Nitrates (NO3-), acetates (CH3COO-), chlorates (ClO3-), and perchlorates (ClO4-) are soluble.

  3. Salts containing Ag+, Pb2+, and Hg22+ are usually insoluble.

  4. Most chlorides (Cl-), bromides (Br-), and iodides (I-) are soluble, except for Ag+, Pb2+, and Hg22+.

  5. Sulfates (SO4^{2-}) are generally soluble except those involving Ca2+, Sr2+, and Ba2+.

  6. Carbonates (CO3^{2-}), hydroxides (OH-), phosphates (PO4^{3-}), and sulfides (S^{2-}) are generally insoluble.

Note for Application:

  • Follow these rules in order!

Net Ionic Equations

  • Types of Chemical Equations:

  1. Molecular Equation: Lists all species in molecular form, e.g.,
    ext{HCl}(aq) + ext{NaOH}(aq)
    ightarrow ext{H}_2 ext{O}(l) + ext{NaCl}(aq)

  2. Complete Ionic Equation: Strong electrolytes are dissociated into ions, e.g.,
    ext{H}^+(aq) + ext{Cl}^-(aq) + ext{Na}^+(aq) + ext{OH}^-(aq)
    ightarrow ext{H}_2 ext{O}(l) + ext{Na}^+(aq) + ext{Cl}^-(aq)

  3. Net Ionic Equation (NIE): Only includes ions that participate in the reaction, excluding spectator ions, e.g.,
    ext{H}^+(aq) + ext{OH}^-(aq)
    ightarrow ext{H}_2 ext{O}(l)

Key Notes:

  1. A compound is represented as ions if it is soluble and a strong electrolyte.

  2. Polyatomic ions remain intact (not split into smaller ions).

  3. Focus only on aqueous species, leaving solids, gases, and liquids alone.

Types of Chemical Reactions

  • Classification of Reactions: Due to the immense variety, reactions can be categorized into several types.

  • Main Types:

    • Precipitation Reactions

    • Acid-base Reactions

    • Oxidation-reduction Reactions

Precipitation Reactions:

  • Definition: A chemical reaction that leads to the formation of an insoluble solid, known as the precipitate.

  • Characteristic: A metathesis reaction where ions are exchanged in solution.

  • General Form:
    ext{AX} + ext{BY}
    ightarrow ext{AY} + ext{BX}

  • Example Reaction:
    2 ext{AgNO}3(aq) + ext{MgCl}2(aq)
    ightarrow 2 ext{AgCl}(s) + ext{Mg(NO}3)2(aq)

Concept Check 4.3

  • Evaluate the following reactions for precipitate formation:

  • Reactions:

  • 1. ext{NiCl}2(aq) + (NH4)_2 ext{S}(aq)
    ightarrow ?

  • 2. ext{Na}2 ext{CrO}4(aq) + ext{Pb(NO}3)2(aq)
    ightarrow ?

  • 3. ext{AgCl}(aq) + ext{CaBr}_2(aq)
    ightarrow ?

  • 4. ext{ZnCl}2(aq) + ext{K}2 ext{CO}_3(aq)
    ightarrow ?

Select the best answer:

  • A. 1

  • B. 2

  • C. 3

  • D. 4