Introduction to Chemical Principles

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  • Edition: 10th
  • Format: Paperback
  • Copyright: 2010-01-04
  • Publisher: Prentice Hall
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List Price: $138.07


This solid, yet value-priced paperback gives you the background and confidence you'll need to succeed in chemistry. Stoker focuses on the most important topics—omitting organic and biochemistry chapters—and teaches the problem-solving skills students in this course need. Each topic is developed at “ground level,” and continues step by step until the level of sophistication required for a further chemistry course is attained.

This interesting and informative book gives readers the background (and confidence) needed for subsequent success in general chemistry. Honing selected portions for greater clarity in presentation, Stoker's book foregoes the multi-topic, “watered down” version of its counterparts by zeroing in on a smaller number of select topics, treating them extensively, and helping readers gain a solid grasp of essential, fundamental material. Develops each topic at “ground level” and continues step-by-step until the level of sophistication required for further chemistry study is attained.

Topics place an emphasis on problem-solving throughout—uses dimensional analysis in problem-solving whenever possible. Features 17 vignettes on “The Human Side of Chemistry”—Brief biographies of scientists who helped develop the foundations of modern chemistry. For those who have had little or no previous instruction in chemistry, or for those who want a thorough review of chemical principles.

Author Biography

Stephen Stoker of Weber State University in Ogden, UT, has taught the gamut of introductory chemistry courses, specializing in GOB, for the past 30 years. Weber State has a very large Health Professions school, and GOB is geared at students with little or no experience in chemistry (notably Allied Health and Nursing majors).

Table of Contents

NOTE: Each chapter concludes with a Summary, Key Terms, Practice 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 and Chemical Technology
1.4 How Chemists Discover Things–The Scientific Method
1.5 The Limitations of the Scientific Method
1.6 The Limitations 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 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 Elements
    THE 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)
4.10 The Atom
    THE HUMAN SIDE OF CHEMISTRY 3: John Dalton (1766—1844)
4.11 The Molecule
4.12 Natural and Synthetic Compounds
4.13 Chemical Formulas

Chapter 5 Subatomic Particles, Isotopes, and Nuclear Chemistry
5.1 Subatomic Particles: Protons, Neutrons, and Electrons
5.2 Atomic Number and Mass Number
5.3 Isotopes
5.4 Atomic Masses
5.5 Evidence Supporting the Existence and Arrangement of Subatomic Particles
5.6 Nuclear Stability and Radioactivity
    THE HUMAN SIDE OF CHEMISTRY 4: Ernest Rutherford (1871—1937)
5.7 Half-Life: A Measure of Nuclear stability
5.8 The Nature of Natural Radioactive Emissions
5.9 Equations for Radioactive Decay
5.10 Transmutation and Bombardment Reactions
5.11 Positron Emission and Electron Capture
5.12 Neutron-to-Proton Ratio and Type of Radioactive Decay
5.13 Radioactive Decay Series

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 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 Nomenclature for Ionic Compounds Containing Polyatomic Ions
8.5 Nomenclature for Binary Molecular Compounds
8.6 Nomenclature for Acids
8.7 Nomenclature Rules–A Summary

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 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 Concentration: Percentage of Solute
13.7 Concentration: Parts per Million and Parts per Billion
13.8 Concentration: Molarity
13.9 Concentration: Molality
13.10 Dilution
13.11 Molarity and Chemical Equations
13.12 Calculations Involving Volume: A Summary

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

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

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