Ch 4 Review of Chemical Reactions & Stoichiometry

Chemical Reactions & Stoichiometry Study Notes

Chapter Overview

  • This chapter focuses on the principles of chemical reactions and stoichiometry as outlined in the textbook by Tro, Fridgen, Shaw.
  • Students are encouraged to read the book and attend lectures for comprehensive understanding.
  • Future chapters will contain less material to help students improve their note-taking skills.

1. Writing & Balancing Chemical Equations

1.1 Introduction

  • Chemical Reaction Definition: A process in which one set of substances (reactants) is converted into another set of substances (products).
    • Evidence for a reaction can be physical or chemical.
    • Mass Balance: For each element, the number of atoms on the reactants side must equal those on the products side.
    • This balances with mass laws and atomic theory.

1.2 Chemical Reaction Equations

  • Chemical Reaction Equation: A symbolic representation showing the identities and quantities of substances involved in a chemical change.
    • Format: Reactants → Products
    • Left side (reactants) indicates the amounts present before the reaction; right side (products) shows the amounts after.
    • Reactants undergo breaking and forming of chemical bonds, observing conservation of mass and total number of atoms.

1.3 Balancing Chemical Equations

  • Example Given: nitrogen monoxide gas + oxygen gas → nitrogen dioxide gas
    1. Translate statement: NO(g) + O2(g) → NO2(g)
    2. Balance atoms:
    • 2 NO(g) + O2(g) → 2 NO2(g)
    • Balancing Strategies:
      • Start with the most complex molecule and end with the least complex.
      • Use balancing by inspection.
      • Treat coefficients as factors for all atoms in the formula.
      • Do not alter chemical formulas; do not add substances.
      • A balanced equation remains valid if multiplied by any whole number.

1.4 Coefficients in Balancing

  • Adjust coefficients while maintaining the smallest whole number ratios.
  • The state symbols for reactants and products are:
    • Solid (s)
    • Liquid (l)
    • Gas (g)
    • Aqueous (aq)

1.5 Example Balancing Problems

  • Example 1: (NH4)2Cr2O7(s) → Cr2O3(s) + N2(g) + H2O(g)

    • Balancing process Indicator:
    • N: 2 (balanced), Cr: 2 (balanced), H: 8 (unbalanced), O: 7 (unbalanced) → Adjust H2O coefficient to 4.

  • Example 2: Combustion of Octane (C8H18)

    • Formulation and balancing steps are shown.
    • Final balanced equation: 2 C8H18(l) + 25 O2(g) → 16 CO2(g) + 18 H2O(g)

1.6 Reaction Conditions

  • When listing reactions, conditions like temperature, pressure, or catalysts may be noted.

2. Solutions & Solubility

2.1 Electrolytes

  • Electrolyte Definition: A substance that dissociates into ions in solution.
    • Strong Electrolytes: Completely dissociate (e.g., NaCl).
    • Weak Electrolytes: Only partially dissociate (e.g., CH3COOH).
    • Non-electrolytes: Do not dissociate (e.g., C12H22O11).

2.2 Dissociation & Ionization

  • Dissociation: Separation of ionic compounds into ions in solution.
  • Ionization: The process of forming ions from molecules (usually in solids).
    • Example: Ionic compounds (solid) produce ions upon dissolution.

2.3 Precipitation Reactions

  • Defined as ionic reactions leading to an insoluble product (precipitate).
  • Important Notes on Solubility:
    • Not all ionic compounds are soluble in water.
    • Solubility rules provide guidance for predicting reactions.
    • For example, all Na+, K+, and NH4+ salts are soluble.
    • Some groups like carbonates and phosphates are generally insoluble unless combined with soluble ions.

2.4 Predicting Outcomes of Precipitation Reactions

  • Determine present ionic species, analyze interactions, and apply solubility rules.
    • Example: AgNO3 (aq) + KCl (aq) → AgCl (s) + KNO3 (aq)

2.5 Types of Chemical Equations in Solutions

  • Molecular Reaction Equation: Shows complete chemical formulas.
  • Complete Ionic Reaction Equation: Separates strong electrolytes into their constituent ions.
  • Net Ionic Reaction Equation: Shows only the species that participate in the reaction (removes spectator ions).

3. Acid-Base Reactions

3.1 Brønsted-Lowry Theory

  • Acid Definition: A substance that produces H+ or H3O+ in solution; a proton donor.
    • Strong acids like HCl dissociate completely in solution.
    • Weak acids such as CH3COOH dissociate incompletely, representing weak electrolytes.

3.2 Base Definition

  • A substance that produces OH– in solution or accepts protons.
    • Strong bases dissociate completely, while weak bases do not.

3.3 Neutralization Reaction

  • The reaction between an acid and base to form water and a salt.

Note: This document is an exhaustive summary of reactions, processes in stoichiometry, acid-base reactions, and other related principles present in the leading chapter on chemical reactions and stoichiometry. Further study should include real applications and worked examples for clarity.