Chemical Principles

Chemical Principles

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Known for helping students develop the qualitative, conceptual foundation that gets them thinking like chemists, this market-leading text is designed for students with solid mathematical preparation and prior exposure to chemistry. The unique organization of the text supports this qualitative-to-quantitative approach. A strong emphasis on models and everyday applications of chemistry combines with a thoughtful, step-by-step problem solving approach to build conceptual understanding.
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Product details

  • Hardback | 1200 pages
  • 208.28 x 256.54 x 45.72mm | 2,154.55g
  • CA, United States
  • English
  • 6th edition
  • 4 colour illustrations
  • 061894690X
  • 9780618946907
  • 2,274,913

Table of contents

Note: Each chapter concludes with Exercises.
1. Chemists and Chemistry
1.1 Thinking Like a Chemist
1.2 A Real-World Chemistry Problem
1.3 The Scientific Method
1.4 Industrial Chemistry
1.5 Polyvinyl Chloride (PVC): Real-World Chemistry
2. Atoms, Molecules, and Ions
2.1 The Early History of Chemistry
2.2 Fundamental Chemical Laws
2.3 Dalton's Atomic Theory
2.4 Cannizzaro's Interpretation
2.5 Early Experiments to Characterize the Atom
2.6 The Modern View of Atomic Structure: An Introduction
2.7 Molecules and Ions
2.8 An Introduction to the Periodic Table
2.9 Naming Simple Compounds
3. Stoichiometry
3.1 Atomic Masses
3.2 The Mole
3.3 Molar Mass
3.4 Percent Composition of Compounds
3.5 Determining the Formula of a Compound
3.6 Chemical Equations
3.7 Balancing Chemical Equations
3.8 Stoichiometric Calculations: Amounts of Reactants and Products
3.9 Calculations Involving a Limiting Reactant
4. Types of Chemical Reactions and Solution Stoichiometry
4.1 Water, the Common Solvent
4.2 The Nature of Aqueous Solutions: Strong and Weak Electrolytes
4.3 The Composition of Solutions
4.4 Types of Chemical Reactions
4.5 Precipitation Reactions
4.6 Describing Reactions in Solution
4.7 Selective Precipitation
4.8 Stoichiometry of Precipitation Reactions
4.9 Acid-Base Reactions
4.10 Oxidation-Reduction Reactions
4.11 Balancing Oxidation-Reduction Equations
4.12 Simple Oxidation-Reduction Titrations
5. Gases
5.1 Early Experiments
5.2 The Gas Laws of Boyle, Charles, and Avogadro
5.3 The Ideal Gas Law
5.4 Gas Stoichiometry
5.5 Dalton's Law of Partial Pressures
5.6 The Kinetic Molecular Theory of Gases
5.7 Effusion and Diffusion
5.8 Collisions of Gas Particles with the Container Walls
5.9 Intermolecular Collisions
5.10 Real Gases
5.11 Characteristics of Several Real Gases
5.12 Chemistry in the Atmosphere
6. Chemical Equilibrium
6.1 The Equilibrium Condition
6.2 The Equilibrium Constant
6.3 Equilibrium Expressions Involving Pressures
6.4 The Concept of Activity
6.5 Heterogeneous Equilibria
6.6 Applications of the Equilibrium Constant
6.7 Solving Equilibrium Problems
6.8 Le Chatelier's Principle
6.9 Equilibria Involving Real Gases
7. Acids and Bases
7.1 The Nature of Acids and Bases
7.2 Acid Strength
7.3 The pH Scale
7.4 Calculating the pH of Strong Acid Solutions
7.5 Calculating the pH of Weak Acid Solutions
7.6 Bases
7.7 Polyprotic Acids
7.8 Acid-Base Properties of Salts
7.9 Acid Solutions in Which Water Contributes to the H+ Concentration
7.10 Strong Acid Solutions in Which Water Contributes to the H+ Concentration
7.11 Strategy for Solving Acid-Base Problems: A Summary
8. Applications of Aqueous Equilibria
8.1 Solutions of Acids or Bases Containing a Common Ion
8.2 Buffered Solutions
8.3 Exact Treatment of Buffered Solutions
8.4 Buffer Capacity
8.5 Titrations and pH Curves
8.6 Acid-Base Indicators
8.7 Titration of Polyprotic Acids
8.8 Solubility Equilibria and the Solubility Product
8.9 Precipitation and Qualitative Analysis
8.10 Complex Ion Equilibria
9. Energy, Enthalpy, and Thermochemistry
9.1 The Nature of Energy
9.2 Enthalpy
9.3 Thermodynamics of Ideal Gases
9.4 Calorimetry
9.5 Hess's Law
9.6 Standard Enthalpies of Formation
9.7 Present Sources of Energy
9.8 New Energy Sources
10. Spontaneity, Entropy, and Free Energy
10.1 Spontaneous Processes and Entropy
10.2 The Isothermal Expansion and Compression of an Ideal Gas
10.3 The Definition of Entropy
10.4 Entropy and Physical Changes
10.5 Entropy and the Second Law of Thermodynamics
10.6 The Effect of Temperature on Spontaneity
10.7 Free Energy
10.8 Entropy Changes in Chemical Reactions
10.9 Free Energy and Chemical Reactions
10.10 The Dependence of Free Energy on Pressure
10.11 Free Energy and Equilibrium
10.12 Free Energy and Work
10.13 Reversible and Irreversible Processes: A Summary
10.14 Adiabatic Processes
11. Electrochemistry
11.1 Galvanic Cells
11.2 Standard Reduction Potentials
11.3 Cell Potential, Electrical Work, and Free Energy
11.4 Dependence of the Cell Potential on Concentration
11.5 Batteries
11.6 Corrosion
11.7 Electrolysis
11.8 Commercial Electrolytic Processes
12. Quantum Mechanics and Atomic Theory
12.1 Electromagnetic Radiation
12.2 The Nature of Matter
12.3 The Atomic Spectrum of Hydrogen
12.4 The Bohr Model
12.5 The Quantum Mechanical Description of the Atom
12.6 The Particle in a Box
12.7 The Wave Equation for the Hydrogen Atom
12.8 The Physical Meaning of a Wave Function
12.9 The Characteristics of Hydrogen Orbitals
12.10 Electron Spin and the Pauli Principle
12.11 Polyelectronic Atoms
12.12 The History of the Periodic Table
12.13 The Aufbau Principle and the Periodic Table
12.14 Further Development of the Polyelectronic Model
12.15 Periodic Trends in Atomic Properties
12.16 The Properties of a Group: The Alkali Metals
13. Bonding: General Concepts
13.1 Types of Chemical Bonds
13.2 Electronegativity
13.3 Bond Polarity and Dipole Moments
13.4 Ions: Electron Configurations and Sizes
13.5 Formation of Binary Ionic Compounds
13.6 Partial Ionic Character of Covalent Bonds
13.7 The Covalent Chemical Bond: A Model
13.8 Covalent Bond Energies and Chemical Reactions
13.9 The Localized Electron Bonding Model
13.10 Lewis Structure
13.11 Resonance
13.12 Exceptions to the Octet Rule
13.13 Molecular Structure: The VSEPR Model
14. Covalent Bonding: Orbitals
14.1 Hybridization and the Localized Electron Model
14.2 The Molecular Orbital Model
14.3 Bonding in Homonuclear Diatomic Molecules
14.4 Bonding in Heteronuclear Diatomic Molecules
14.5 Combining the Localized Electron and Molecular Orbital Models
14.6 Orbitals: Human Inventions
14.7 Molecular Spectroscopy: An Introduction
14.8 Electronic Spectroscopy
14.9 Vibrational Spectroscopy
14.10 Rotational Spectroscopy
14.11 Nuclear Magnetic Resonance Spectroscopy
15. Chemical Kinetics
15.1 Reaction Rates
15.2 Rate Laws: An Introduction
15.3 Determining the Form of the Rate Law
15.4 The Integrated Rate Law
15.5 Rate Laws: A Summary
15.6 Reaction Mechanisms
15.7 The Steady-State Approximation
15.8 A Model for Chemical Kinetics
15.9 Catalysis
16. Liquids and Solids
16.1 Intermolecular Forces
16.2 The Liquid State
16.3 An Introduction to Structures and Types of Solids
16.4 Structure and Bonding in Metals
16.5 Carbon and Silicon: Network Atomic Solids
16.6 Molecular Solids
16.7 Ionic Solids
16.8 Structures of Actual Ionic Solids
16.9 Lattice Defects
16.10 Vapor Pressure and Changes of State
16.11 Phase Diagrams
16.12 Nanotechnology
17. Properties of Solutions
17.1 Solution Composition
17.2 The Thermodynamics of Solution Formation
17.3 Factors Affecting Solubility
17.4 The Vapor Pressures of Solutions
17.5 Boiling-Point Elevation and Freezing-Point Depression
17.6 Osmotic Pressure
17.7 Colligative Properties of Electrolyte Solutions
17.8 Colloids
18. The Representative Elements
18.1 A Survey of the Representative Elements
18.2 The Group 1A Elements
18.3 The Chemistry of Hydrogen
18.4 The Group 2A Elements
18.5 The Group 3A Elements
18.6 The Group 4A Elements
18.7 The Group 5A Elements
18.8 The Chemistry of Nitrogen
18.9 The Chemistry of Phosphorus
18.10 The Group 6A Elements
18.11 The Chemistry of Oxygen
18.12 The Chemistry of Sulfur
18.13 The Group 7A Elements
18.14 The Group 8A Elements
19. Transition Metals and Coordination Chemistry
19.1 The Transition Metals: A Survey
19.2 The First-Row Transition Metals
19.3 Coordination Compounds
19.4 Isomerism
19.5 Bonding in Complex Ions: The Localized Electron Model
19.6 The Crystal Field Model
19.7 The Molecular Orbital Model
19.8 The Biological Importance of Coordination Complexes
20. The Nucleus: A Chemist's View
20.1 Nuclear Stability and Radioactive Decay
20.2 The Kinetics of Radioactive Decay
20.3 Nuclear Transformations
20.4 Detection and Uses of Radioactivity
20.5 Thermodynamic Stability of the Nucleus
20.6 Nuclear Fission and Nuclear Fusion
20.7 Effects of Radiation
21. Organic and Biochemical Molecules
21.1 Alkanes: Saturated Hydrocarbons
21.2 Alkenes and Alkynes
21.3 Aromatic Hydrocarbons
21.4 Hydrocarbon Derivatives
21.5 Polymers
21.6 Natural Polymers
Appendix 1. Mathematical Procedures
A1.1 Exponential Notation
A1.2 Logarithms
A1.3 Graphing Functions
A1.4 Solving Quadratic Equations
A1.5 Uncertainties in Measurements
A1.6 Significant Figures
Appendix 2. Units of Measurement and Conversions Among Units
A2.1 Measurements
A2.2 Unit Conversions
Appendix 3. Spectral Analysis
Appendix 4. Selected Thermodynamic Data
Appendix 5. Equilibrium Constants and Reduction Potentials
Answers to Selected Exercises
Photo Credits
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Review quote

My sense is that the topics are pitched at an appropriate level. The writing style is crisp and to the point whereas the art, tables, and end of chapter problems nicely support the written text. The writing style is lucid, direct, precise, and non-verbose. Some other introductory chemistry texts adopt a verbose presentation with the attendant loss of clarity. The students who are floundering and to whom I have recommended Zumdahl invariably come away with an enhanced understanding of the relevant chemical concepts after a thorough encounter with this text. This is the singular text that does not compromise discipline of mind and contains flashes of insight and brilliance reminiscent of the Linus Pauling series of introductory texts, and yet is conceptually and algebraically accessible to most students with a good background in high school chemistry.
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About Steven Zumdahl

Steve Zumdahl is the author of market-leading textbooks in introductory chemistry, general chemistry, honors-level chemistry, and high school chemistry. Recently retired from his long-time position as Director of Undergraduate Programs at the University of Illinois, he has received numerous awards for his contributions to chemical education. These include the National Catalyst Award in recognition of his contribution to chemical education, the University of Illinois Teaching Award, the UIUC Liberal Arts and Sciences Advising Award, and the School of Chemical Sciences Teaching Award (five times). He earned his B.S. in Chemistry from Wheaton College (IL), and his Ph.D. from the University of Illinois.
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Rating details

74 ratings
3.43 out of 5 stars
5 26% (19)
4 26% (19)
3 26% (19)
2 12% (9)
1 11% (8)
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