ch08
Chapter Overview
Title: Chemical Reactions and Aqueous Solutions
Source: Interactive General Chemistry 2.0, © 2023 Macmillan Learning
Chapter Outline
8.1: Chemical Equations
8.2: Types of Chemical Reactions
8.3: Compounds in Aqueous Solution
8.4: Precipitation Reactions
8.5: Acid-Base Reactions
8.6: Oxidation States and Redox Reactions
8.7: Predicting the Products of Redox Reactions
Section 8.1: Chemical Equations
Balanced Chemical Equations: Show reactants and products and must adhere to the law of conservation of mass.
Phases in Chemical Equations: Indicate physical states of reactants/products (e.g., (s), (l), (g), (aq)).
Example equation: 2 H₂ + O₂ → 2 H₂O (products are in liquid state).
Coefficients: Represent the proportions of reactants/products in a balanced reaction.
Information in a Chemical Equation
Reactants Rearrangement: Reactants rearrange to form products.
Example: 2 H₂ + O₂ → 2 H₂O (hydrogens and oxygen rearrange to form water).
Physical States Notation: Included after each formula.
Example: 1 mol MgI2(aq) + 2 mol AgNO3(aq) → 2 mol AgI(s) + 1 mol Mg(NO3)2(aq)
Reaction Conditions
Noted above or below the reaction arrow (e.g., heat application for reactions to proceed).
Balancing Equations
Law of Conservation of Mass: All atoms from reactants must equal the atoms in products.
Steps for Balancing:
Change coefficients, never subscripts.
Balance polyatomic ions as units if both reactants and products.
Keep a balance order (left to right).
Final check for all atoms and smallest whole-number ratios.
Example Balancing Exercises
Example 8.1: Balancing Ba(OH)2 + HBr → BaBr2 + H2O.
Rationale and verification of atom counts post-balancing.
Section 8.2: Types of Chemical Reactions
Combination Reactions: Two simple compounds combine to form a complex product.
Example: 4 Fe + 3 O2 → 2 Fe2O3 (rust formation).
Decomposition Reactions: A single reactant breaks down into less complex products.
Example: 2 H2O → 2 H2 + O2 (electrolysis).
Single-Replacement Reactions: An element displaces another in a compound.
Example: Zn + 2 HCl → ZnCl2 + H2.
Double-Replacement Reactions: Ionic compounds exchange ions to form new compounds.
Example: KI + Pb(NO3)2 → PbI2 + KNO3.
Acid-Base Reactions: A type of double-replacement where an acid and base react to form salt and water.
Combustion Reactions: A substance combines rapidly with oxygen, producing heat and light.
Example: CH₄ + 2 O₂ → CO₂ + 2 H₂O.
Driving Forces for Reactions in Aqueous Solutions
Spontaneous reactions lead to stable, lower energy products; involves changes in enthalpy and entropy.
Driving Forces of Types:
Precipitation reactions produce insoluble compounds.
Neutralization reactions (acid-base) produce salts and water.
Section 8.3: Compounds in Aqueous Solution
Solubility Definitions:
Soluble: Dissolves in water; Insoluble: Remains as solid.
Electrolytes: Compounds that conduct electricity when dissolved.
Strong Electrolytes: Dissociate completely (e.g., NaCl).
Weak Electrolytes: Partially dissociate (e.g., acetic acid).
Section 8.4: Precipitation Reactions
Precipitation Prediction: When two ionic compounds form an insoluble product in solution.
Use solubility guidelines to predict and identify precipitates.
Section 8.5: Acid-Base Reactions
General Reaction: Acid + Base → Salt + Water.
Net Ionic Equations: Show only species involved in the reaction excluding spectator ions.
Section 8.6: Oxidation States and Redox Reactions
Oxidation States: Assist in tracking electron transfer during reactions. Rules define how to assign oxidation states.
Redox Reactions: Define changes in oxidation states indicating oxidation (loss of electrons) and reduction (gain of electrons).
Example (oxidation): 4Al + 3O2 → 2Al2O3 (Al oxidized, O reduced).
Section 8.7: Predicting the Products of Redox Reactions
Use activity series to determine reactions and predict products.
Single-Replacement Reactions need comparison of reactivity of metals (element displaces another).