Introduction to Chemical Principles, 11th edition

Table of Contents

NOTE: Each chapter concludes with Concepts to Remember, Key Terms, Practice Problems, Multi-Concept Problems and Multiple-Choice Practice Test.

• Chapter 1: The Science of Chemistry
  • 1.1 Chemistry-A Scientific Discipline
  • 1.2 Scientific Research and Technology
  • 1.3 The Scope of Chemistry
  • 1.4 How Chemists Discover Things-The Scientific Method
  • 1.5 The Limitations of the Scientific Method
  • 1.6 Application Limitations for Methods of Science
• Chapter 2: Numbers from Measurements
  • 2.1 The Importance of Measurement
  • 2.2 Exact and Inexact Numbers
  • 2.3 Accuracy, Precision, and Error
  • 2.4 Uncertainty in Measurements
  • 2.5 Significant Figures
  • 2.6 Significant Figures and Mathematical Operations
  • 2.7 Expressing Numbers in Scientific Notation
  • 2.8 Mathematical Operations in Scientific Notation
• Chapter 3: Unit Systems and Dimensional Analysis
  • 3.1 The Metric System of Units
  • 3.2 Metric Units of Length
  • 3.3 Metric Units of Mass
  • 3.4 Metric Units of Volume
  • 3.5 Units in Mathematical Operations
  • 3.6 Conversion Factors
  • 3.7 Dimensional Analysis
  • 3.8 Density
  • 3.9 Equivalence Conversion Factors Other Than Density
  • 3.10 Percentage and Percent Error
  • 3.11 Temperature Scales
• Chapter 4: Basic Concepts About Matter
  • 4.1 Chemistry-The Study of Matter
  • 4.2 Physical States of Matter
  • 4.3 Properties of Matter
  • 4.4 Changes in Matter
  • 4.5 Pure Substances and Mixtures
  • 4.6 Heterogeneous and Homogeneous Mixtures
  • 4.7 Elements and Compounds
  • 4.8 Discovery and Abundance of the ElementsTHE HUMAN SIDE OF CHEMISTRY 1: Joseph Priestley (1733–1804)
  • 4.9 Names and Chemical Symbols of the Elements
    • THE HUMAN SIDE OF CHEMISTRY 2: Jöns Jakob Berzelius (1779–1848)
• Chapter 5: Atoms, Molecules, and Subatomic Particles
  • 5.1 The Atom
    • THE HUMAN SIDE OF CHEMISTRY 3: John Dalton (1766–1844)
  • 5.2 The Molecule
  • 5.3 Natural and Synthetic Compounds
  • 5.4 Chemical Formulas
  • 5.5 Subatomic Particles: Protons, Neutrons, and Electrons
  • 5.6 Atomic Number and Mass Number
  • 5.7 Isotopes
  • 5.8 Atomic Masses
  • 5.9 Evidence Supporting the Existence and Arrangement of Subatomic Particles
    • THE HUMAN SIDE OF CHEMISTRY 4: Ernest Rutherford (1871–1937)
• Chapter 6: Electronic Structure and Chemical Periodicity
  • 6.1 The Periodic Law
  • 6.2 The Periodic Table
    • THE HUMAN SIDE OF CHEMISTRY 5: Dmitri Ivanovich Mendeleev (1834–1907)
  • 6.3 The Energy of an Electron
    • THE HUMAN SIDE OF CHEMISTRY 6: Erwin Schrödinger (1887–1961)
  • 6.4 Electron Shells
  • 6.5 Electron Subshells
  • 6.6 Electron Orbitals
  • 6.7 Electron Configurations
  • 6.8 Electron Orbital Diagrams
  • 6.9 Electron Configurations and the Periodic Law
  • 6.10 Electron Configurations and the Periodic Table
  • 6.11 Classification Systems for the Elements
  • 6.12 Chemical Periodicity
• Chapter 7: Chemical Bonds
  • 7.1 Types of Chemical Bonds
  • 7.2 Valence Electrons and Lewis Symbols
    • THE HUMAN SIDE OF CHEMISTRY 7: Gilbert Newton Lewis (1875–1946)
  • 7.3 The Octet Rule
  • 7.4 The Ionic Bond Model
  • 7.5 The Sign and Magnitude of Ionic Charge
  • 7.6 Lewis Structures for Ionic Compounds
  • 7.7 Chemical Formulas for Ionic Compounds
  • 7.8 Structure of Ionic Compounds
  • 7.9 Polyatomic Ions
  • 7.10 The Covalent Bond Model
  • 7.11 Lewis Structures for Molecular Compounds
  • 7.12 Single, Double, and Triple Covalent Bonds
  • 7.13 Valence Electron Count and Number of Covalent Bonds Formed
  • 7.14 Coordinate Covalent Bonds
  • 7.15 Resonance Structures
  • 7.16 Systematic Procedures for Drawing Lewis Structures
  • 7.17 Molecular Geometry
  • 7.18 Electronegativity
    • THE HUMAN SIDE OF CHEMISTRY 8: Linus Carl Pauling (1901–1994)
  • 7.19 Bond Polarity
  • 7.20 Molecular Polarity
• Chapter 8: Chemical Nomenclature
  • 8.1 Classification of Compounds for Nomenclature Purposes
  • 8.2 Types of Binary Ionic Compounds
  • 8.3 Nomenclature for Binary Ionic Compounds
  • 8.4 Chemical Formulas for Polyatomic Ions
  • 8.5 Nomenclature for Ionic Compounds Containing Polyatomic Ions
  • 8.6 Nomenclature for Binary Molecular Compounds
  • 8.7 Nomenclature for Acids
  • 8.8 System Procedures for Using Nomenclature Rules
• Chapter 9: Chemical Calculations: The Mole Concept and Chemical Formulas
  • 9.1 The Law of Definite Proportions
    • THE HUMAN SIDE OF CHEMISTRY 9: Joseph-Louis Proust (1754–1826)
  • 9.2 Calculation of Formula Masses
  • 9.3 Significant Figures and Atomic Mass
  • 9.4 Mass Percent Composition of a Compound
  • 9.5 The Mole: The Chemist's Counting Unit
    • THE HUMAN SIDE OF CHEMISTRY 10: Lorenzo Romano Amedeo Carlo Avogadro (1776–1856)
  • 9.6 The Mass of a Mole
  • 9.7 Significant Figures and Avogadro's Number
  • 9.8 Relationship between Atomic Mass Units and Gram Units
  • 9.9 The Mole and Chemical Formulas
  • 9.10 The Mole and Chemical Calculations
  • 9.11 Purity of Samples
  • 9.12 Empirical and Molecular Formulas
  • 9.13 Determination of Empirical Formulas
  • 9.14 Determination of Molecular Formulas
• Chapter 10: Chemical Calculations Involving Chemical Equations
  • 10.1 The Law of Conservation of Mass
    • THE HUMAN SIDE OF CHEMISTRY 11: Antoine-Laurent Lavoisier (1743–1794)
  • 10.2 Writing Chemical Equations
  • 10.3 Chemical Equation Coefficients
  • 10.4 Balancing Procedures for Chemical Equations
  • 10.5 Special Symbols Used in Chemical Equations
  • 10.6 Classes of Chemical Reactions
  • 10.7 Chemical Equations and the Mole Concept
  • 10.8 Balanced Chemical Equations and the Law of Conservation of Mass
  • 10.9 Calculations Based on Chemical Equations-Stoichiometry
  • 10.10 The Limiting Reactant Concept
  • 10.11 Yields: Theoretical, Actual, and Percent
  • 10.12 Simultaneous and Sequential Chemical Reactions
• Chapter 11: States of Matter
  • 11.1 Factors That Determine Physical State
  • 11.2 Property Differences among Physical States
  • 11.3 The Kinetic Molecular Theory of Matter
  • 11.4 The Solid State
  • 11.5 The Liquid State
  • 11.6 The Gaseous State
  • 11.7 A Comparison of Solids, Liquids, and Gases
  • 11.8 Endothermic and Exothermic Changes of State
  • 11.9 Heat Energy and Specific Heat
  • 11.10 Temperature Changes as a Substance Is Heated
  • 11.11 Energy and Changes of State
  • 11.12 Heat Energy Calculations
  • 11.13 Evaporation of Liquids
  • 11.14 Vapor Pressure of Liquids
  • 11.15 Boiling and Boiling Points
  • 11.16 Intermolecular Forces in Liquids
  • 11.17 Hydrogen Bonding and the Properties of Water
• Chapter 12: Gas Laws
  • 12.1 Properties of Some Common Gases
  • 12.2 Gas Law Variables
  • 12.3 Boyle's Law: A Pressure—Volume Relationship
    • THE HUMAN SIDE OF CHEMISTRY 12: Robert Boyle (1627–1691)
  • 12.4 Charles's Law: A Temperature—Volume Relationship
    • THE HUMAN SIDE OF CHEMISTRY 13: Jacques Alexandre César Charles (1746–1823)
  • 12.5 Gay-Lussac's Law: A Temperature—Pressure Relationship
    • THE HUMAN SIDE OF CHEMISTRY 14: Joseph Louis Gay-Lussac (1778–1850)
  • 12.6 The Combined Gas Law
  • 12.7 Avogadro's Law
  • 12.8 An Ideal Gas
  • 12.9 The Ideal Gas Law
  • 12.10 Modified Forms of the Ideal Gas Law Equation
  • 12.11 Volumes of Gases in Chemical Reactions
  • 12.12 Volumes of Gases and the Limiting Reactant Concept
  • 12.13 Molar Volume of a Gas
  • 12.14 Chemical Calculations Using Molar Volume
  • 12.15 Mixtures of Gases
  • 12.16 Dalton's Law of Partial Pressures
• Chapter 13: Solutions
  • 13.1 Characteristics of Solutions
  • 13.2 Solubility
  • 13.3 Solution Formation
  • 13.4 Solubility Rules
  • 13.5 Solution Concentrations
  • 13.6 Percentage Concentration Unit
  • 13.7 Parts per Million and Parts per Billion Concentration Unit
  • 13.8 Molarity Concentration Units
  • 13.9 Molality and Chemical Reactions in Solution
  • 13.10 Dilution Calculations
  • 13.11 Molarity Concentration Unit
• Chapter 14: Acids, Bases, and Salts
  • 14.1 Arrhenius Acid—Base Theory
    • THE HUMAN SIDE OF CHEMISTRY 15: Svante August Arrhenius (1859–1927)
  • 14.2 Brønsted—Lowry Acid—Base Theory
  • 14.3 Conjugate Acids and Bases
  • 14.4 Mono-, Di-, and Triprotic Acids
  • 14.5 Strengths of Acids and Bases
  • 14.6 Salts
  • 14.7 Reactions of Acids
  • 14.8 Reactions of Bases
  • 14.9 Reactions of Salts
  • 14.10 Self-Ionization of Water
  • 14.11 The pH Scale
  • 14.12 Hydrolysis of Salts
  • 14.13 Buffers
  • 14.14 Acid-Base Titrations
• Chapter 15: Chemical Equations: Net Ionic and Oxidation-Reduction
  • 15.1 Types of Chemical Equations
  • 15.2 Electrolytes
  • 15.3 Ionic and Net Ionic Equations
  • 15.4 Oxidation—Reduction Terminology
  • 15.5 Oxidation Numbers
  • 15.6 Redox and Nonredox Chemical Reactions
  • 15.7 Balancing Oxidation—Reduction Equations
  • 15.8 Oxidation Number Method for Balancing Redox Equations
  • 15.9 Half-Reaction Method for Balancing Redox Equations
  • 15.10 Disproportionation Reactions
  • 15.11 Stoichiometric Calculations Involving Ions
• Chapter 16: Reaction Rates and Chemical Equilibrium
  • 16.1 Collision Theory
  • 16.2 Endothermic and Exothermic Chemical Reactions
  • 16.3 Factors That Influence Chemical Reaction Rates
  • 16.4 Chemical Equilibrium
  • 16.5 Equilibrium Mixture Stoichiometry
  • 16.6 Equilibrium Constants
  • 16.7 Equilibrium Position
  • 16.8 Temperature Dependency of Equilibrium Constants
  • 16.9 Le Châtelier's Principle
    • THE HUMAN SIDE OF CHEMISTRY 16: Henri-Louis Le Châtelier (1850–1936)
  • 16.10 Forcing Chemical Reactions to Completion

Glossary

Answer to Odd-Numbered Problems and All Self-Test Problems

Index