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Introduction to Chemical Principles, 11th edition

  • H Stephen Stoker

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Overview

Newly updated based on extensive reviewer feedback, this introductory text remains focused on the essentials necessary for success in General Chemistry. Introduction to Chemistry Principles, Eleventh Edition focuses on the most important topics – omitting organic and biochemistry chapters – and teaches the problem-solving skills students need. Each topic is introduced and developed step by step until reaching the level of sophistication required for further course work. This two-color paperback is available through the Pearson Custom Library, giving you the flexibility to choose course content while controlling the cost of the text to your student.

Published by Pearson (July 14th 2021) - Copyright © 2014

ISBN-13: 9780137524181

Subject: Chemistry

Category: Introduction to Chemistry

Overview

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

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