General Chemistry

General Chemistry

by Linus Pauling


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General Chemistry by Linus Pauling

"An excellent text, highly recommended." — Choice
When it was first published, this first-year chemistry text revolutionized the teaching of chemistry by presenting it in terms of unifying principles instead of as a body of unrelated facts. Those principles included modern theories of atomic and molecular structure, quantum mechanics, statistical mechanics, and thermodynamics. In addition, Dr. Pauling attempted to correlate the theories with descriptive chemistry, the observed properties of substances, to introduce the student to the multitude of chemical substances and their properties.
In this extensively revised and updated third edition, the Nobel Prize–winning author maintains an excellent balance between theoretical and descriptive material, although the amount of descriptive chemistry has been decreased somewhat, and the presentation of the subject, especially in relation to the nonmetals, has been revised in such a way as to permit greater correlation with the electronic structure of atoms, especially electronegativity.
The principles of quantum mechanics are discussed on the basis of the de Broglie wavelength of the electron. The quantized energy levels of a particle in a box are derived by means of a simple assumption about the relation of the de Broglie waves to the walls of the box. No attempt is made to solve the Schrödinger wave equation for other systems, but the wave functions of hydrogen-like electrons are presented and discussed in some detail, and the quantum states for other systems are also covered. Statistical mechanics is introduced before thermodynamics, and the discussion of thermodynamics is based on it. This arrangement reflects the author's belief that beginning students can understand statistical mechanics better than chemical thermodynamics.
Aimed at first-year college students who plan to major in chemistry or closely related fields, the book is written in a logical, clear, and understandable style. In addition, many excellent figures are included, along with numerous problems and 75 pages of appendices covering such topics as symmetry of molecules and crystals, hybrid bond orbitals, and magnetic properties of substances.

Product Details

ISBN-13: 9780486656229
Publisher: Dover Publications
Publication date: 04/01/1988
Series: Dover Books on Chemistry
Pages: 992
Sales rank: 314,305
Product dimensions: 5.38(w) x 8.50(h) x (d)

About the Author

Linus Pauling: Two-Time Nobel Laureate
In 1985 Dover reprinted Introduction to Quantum Mechanics with Applications to Chemistry, a well-known older book by Linus Pauling and E. Bright Wilson. This book had been first published fifty years earlier and remarkably still found readers in 1985, and still does today, twenty-five years further on.

The first edition of Pauling's General Chemistry was a short book of less than 250 pages published in 1944, during World War II. Three years later, it had more than doubled in size to almost 600 pages, and the 1953 edition was over 700 pages. Fifteen years later, for the 1970 edition, it reached its final size and configuration at almost 1,000 pages ― and that is the edition which Dover reprinted in 1988. Dr. Pauling's one request at that time was that we keep the price affordable for students.

Linus Pauling is of course the only Dover author to win two Nobel prizes, for Chemistry in 1954 and for Peace in 1962; he is the only winner in history of two unshared Nobel Prizes.

In the Author's Own Words:
"Satisfaction of one's curiosity is one of the greatest sources of happiness in life."

"Do unto others 20% better than you would expect them to do unto you, to correct for subjective error."

"The way to get good ideas is to get lots of ideas, and throw the bad ones away."

"Facts are the air of scientists. Without them you can never fly." — Linus Pauling

Critical Acclaim for General Chemistry:
"An excellent text, highly recommended." — Choice

Read an Excerpt

General Chemistry

By Linus Pauling

Dover Publications, Inc.

Copyright © 1970 Linus Pauling
All rights reserved.
ISBN: 978-0-486-13465-9


The Nature and Properties of Matter

1-1. Matter and Chemistry

The universe is composed of matter and radiant energy. Matter (from the Latin materia, meaning wood or other material) may be defined as any kind of mass-energy (see Section 1-2) that moves with velocities less than the velocity of light, and radiant energy as any kind of mass-energy that moves with the velocity of light.

The different kinds of matter are called substances. Chemistry is the science of substances—their structure, their properties, and the reactions that change them into other substances.

This definition of chemistry is both too narrow and too broad. It is too narrow because the chemist in his study of substances must also study radiant energy, in its interaction with substances. He may be interested in the color of substances, which is produced by the absorption of light. Or he may be interested in the atomic structure of substances, as determined by the diffraction of x-rays (Section 3-7 and Appendix IV) or by the absorption or emission of radiowaves by the substances.

On the other hand, the definition is too broad, in that almost all of science could be included within it. The astrophysicist is interested in the substances that are present in stars and other celestial bodies, or that are distributed, in very low concentration, through interstellar space. The nuclear physicist studies the substances that constitute the nuclei of atoms. The biologist is interested in the substances that are present in living organisms. The geologist is interested in the substances, called minerals, that make up the earth. It is hard to draw a line between chemistry and other sciences.

1-2. Mass and Energy

Matter has mass, and any portion of matter on the earth is attracted toward the center of the earth by the force of gravity; this attraction is called the weight of the portion of matter. For many years scientists thought that matter and radiant energy could be distinguished through the possession of mass by matter and the lack of possession of mass by energy. Then, early in the present century (1905), it was pointed out by Albert Einstein (1879-1955) that energy also has mass, and that light is accordingly attracted by matter through gravitation. This was verified by astronomers, who found that a ray of light traveling from a distant star to the earth and passing close by the sun is bent toward the sun by its gravitational attraction. The observation of this phenomenon was made during a solar eclipse, when the image of the star could be seen close to the sun.

The amount of mass associated with a definite amount of energy is given by an important equation, the Einstein equation, which is an essential part of the theory of relativity:

E = mc2 (1-1)

In this equation E is the amount of energy (J), m is the mass (kg), and c is the velocity of light (m s-1). The velocity of light, c, is one of the fundamental constants of nature; its value is 2.9979 × 108 meters per second.

Until the present century it was also thought that matter could not be created or destroyed, but could only be converted from one form into another. In recent years it has, however, been found possible to convert matter into radiant energy, and to convert radiant energy into matter. The mass m of the matter obtained by the conversion of an amount E of radiant energy or convertible into this amount of radiant energy is given by the Einstein equation. Experimental verification of the Einstein equation has been obtained by the study of processes involving nuclei of atoms. The nature of these processes will be described in later chapters in this book.

Until early in the present century scientists made use of a law of conservation of matter and a law of conservation of energy. These two conservation laws were then combined into a single one, the law of conservation of mass, in which the mass to be conserved includes both the mass of the matter in the system and the mass of the radiant energy in the system.

1-3. The International System of Units

The metric system of units of length, mass, force, and other physical quantities was developed during the French Revolution. Because of their greater convenience and simplicity, metric units have replaced native units (such as the foot and the pound) in scientific work everywhere and have been formally accepted for practical use in many countries (all except the United States, Canada, and some African countries). An extended and improved form of the metric system, called the International System (IS, or sometimes SI, for Systéme International ), was formally adopted by the General Conference of Weights and Measures in 1960.

The symbols of the basic IS units and of the prefixes for fractions and multiples and those for some derived IS units are given in Appendix I. If you have made use of the MKS system (meter-kilogram-second system) in your study of physics the IS system will be familiar to you, for the most part, but if you have made use of the cgs system (centimeter-gram-second system) you will have to learn some new units.

The IS unit of mass, the kilogram, is defined as the mass of a standard object made of a platinum-iridium alloy and kept in Paris. One pound is equal approximately to 453.59 g, and hence 1 kg is equal approximately to 2.205 lb. (Note that it has become customary for the abbreviation of units in the metric system to be written without periods.) There is at the present time a flaw in the International System, in that the name for the unit of mass involves a prefix, kilo. This flaw will remain until agreement about a new name and symbol has been reached. In the meantime we must remember that 1 milligram (symbol 1 mg, not 1 ?kg) is one millionth of the unit of mass, not one thousandth, as indicated by the prefix milli.

The IS unit of length, the meter (m), is equal to about 39.37 inches (1 inch equals exactly 2.54 cm). The meter was formerly defined as the distance between two engraved lines on a standard platinum-iridium bar kept in Paris by the International Bureau of Weights and Measures; in 1960 it was redefined, by international agreement, as 1,650,763.73 wavelengths of the orange-red spectral line of krypton 86.

The IS unit of time is the second (s). It is defined as the interval occupied by 9,192,631,770 cycles of the microwave line of cesium 133 with wavelength about 3.26 cm. The second was formerly defined as 1/86400th of the mean solar day.

The IS unit of volume is the cubic meter, m3. In chemistry a unit that is much used is the liter, symbol 1, which is 1 × 103 m3. The milliliter, 1 × 10-3 1, is equal to the cubic centimeter: 1 ml = 1 cm3.

The IS unit of force is the newton (N), which is defined as the force needed to accelerate a mass of 1 kg by 1 m s-2. The newton is 105 dyne (the dyne, the unit of force in the cgs system, is the force that accelerates 1 g by 1 cm s-2). The IS unit of energy, the joule (J), is the work done by 1 newton in the distance 1 meter: 1 J = 1 N m = 107 erg = 107 dyne cm.

In chemistry the calorie has been extensively used as the unit of energy. The thermochemical calorie, defined as 4.184 J (Appendix I), is approximately the amount of energy needed to raise the temperature of 1 g of water by 1°C. The large calorie (kcal or Cal) is 103 cal. In this book we shall use the joule in most of the tables and discussions. Since most thermochemical reference books use the calorie or kilocalorie, you will find it worth while to remember the conversion factor:

1 cal = 4.184 J

1 kcal = 1 Cal = 4.184 kJ

Example 1-1. Niagara Falls (Horseshoe) is 160 feet high. How much warmer is the water at the bottom than at the top, as the result of the conversion of potential energy into thermal energy? The standard acceleration of gravity is 9.80665 m s-2.

Solution. The gravitational force on a mass of 1 kg at the earth's surface is 9.80665 N. The change is potential energy of 1 kg over a vertical distance h (in meters) is 9.80665 × h J. In this problem h has the value 0.3048 × 160 = 48.77 m (conversion factor from Appendix I); hence the change in potential energy produces 9.80665 × 48.77 = 478 J of thermal energy. The energy required to raise the temperature of 1 kg of water by 1°C is given above as 1 kcal = 4.184 kJ = 4184 J. Hence the increase in temperature of the water is 478/4184 = 0.114°C.

Example 1-2. When 2 kg of uranium 235 undergoes nuclear fission (as in the detonation of the Hiroshima atomic bomb on 6 August 1945), 1.646 × 1014 J of radiant energy and thermal energy is liberated. What is the mass of the material products of the reaction?

Solution. We can calculate the mass of the liberated energy by the use of the Einstein equation (1-1). Rewriting this equation by dividing each side by c2 and introducing the values of E and c, we obtain

m = E/c2 = 1.646 × 1014J/(2.998 × 108)2 m2s-2 = 0.183 × 10-2 kg

Thus, the material mass of 2 kg has decreased by 0.00183 kg (that is, by 0.0915%), leaving material products of the reaction with mass 1.99817 kg.

The Einstein relation between mass and energy has been verified by the direct measurement of the mass of the products and of the energy emitted in nuclear reactions.

Example 1-3. It is found by experiment that when 1 kg of glyceryl trinitrate (nitroglycerine) is exploded, the amount 8.0 × 106 J of energy is liberated. What is the mass of the products of the explosion?

Solution. This example is to be solved in exactly the same way as the preceding one. The mass of the radiant energy that is produced by the explosion is obtained by dividing the energy, E, by the square of the velocity of light:

m = E/c2 = 8.0 × 106J/(2.998 × 108)2 m2 s-2 = 0.89 × 10-10 kg

Thus we calculate that the mass of the products of the explosion is 0.999999999911 kg.

We see that the mass of the products of this chemical reaction differs very slightly from the mass of the reactant—so slightly that it is impossible to detect the change in a direct way. The change, less than one part in ten billion (1 in 1010), is so small that for practical purposes we may say that there is conservation of mass in ordinary chemical reactions.

1-4. Temperature

If two objects are placed in contact with one another, thermal energy may flow from one object to the other one. Temperature is the quality that determines the direction in which thermal energy flows—it flows from the object at higher temperature to the object at lower temperature.

Temperatures are ordinarily measured by means of a thermometer, such as the ordinary mercury thermometer, consisting of a quantity of mercury in a glass tube. The temperature scale used by scientists is the centigrade or Celsius scale; it was introduced by Anders Celsius (1701-1744), a Swedish professor of astronomy, in 1742. On this scale the temperature of freezing water saturated with air is 0°C and the temperature of boiling water is 100°C at 1 atm pressure.

On the Fahrenheit scale, used in everyday life in English-speaking countries, the freezing point of water is 32°F and the boiling point of water is 212°F. On this scale the freezing point and the boiling point differ by 180°, rather than the 100° of the centigrade scale.

To convert temperatures from one scale to another, you need only remember that the Fahrenheit degree is 100/180 or 5/9 of the centigrade degree, and that 0°C is the same temperature as 32°F.

The Kelvin Temperature Scale

About 200 years ago scientists noticed that a sample of gas that is cooled decreases in volume in a regular way, and they saw that if the volume were to continue to decrease in the same way it would become zero at about—273°C. The concept was developed that this temperature,—273°C (more accurately,—273.15°C), is the minimum temperature, the absolute zero. A new temperature scale was then devised by Lord Kelvin, a great British physicist (1824-1907). The Kelvin scale is defined in such a way as to permit the laws of thermodynamics to be expressed in simple form (see Chapter 10).

The IS temperature scale is the Kelvin scale with a new definition of the degree. The absolute zero is taken to be 0°K and the triple point of water is taken to be 273.16°K. (The triple point of water, the temperature at which pure liquid water, ice, and water vapor are in equilibrium, is discussed in Section 11-9.) With this definition of the degree, the boiling point of water at one atmosphere pressure is 373.15°K and the freezing point of water saturated with air at one atmosphere pressure is 273.15°K. Hence the IS Kelvin temperature is 273.15°K greater than the centigrade temperature.

1-5. Kinds of Matter

We shall first distinguish between objects and kinds of matter. An object, such as a human being, a table, a brass doorknob, may be made of one kind of matter or of several kinds of matter. The chemist is primarily interested not in the objects themselves, but in the kinds of matter of which they are composed. He is interested in the alloy brass, whether it is in a doorknob or in some other object; and his interest may be primarily in those properties of the material that are independent of the nature of the objects containing it.


The word material is used in referring to any kind of matter, whether homogeneous or heterogeneous.

A heterogeneous material is a material that consists of parts with different properties. A homogeneous material has the same properties throughout.

Wood, with soft and hard rings alternating, is obviously a heterogeneous material, as is also granite, in which grains of three different substances (the minerals quartz, mica, and feldspar) can be seen.

A mineral is any chemical element, compound, or other homogeneous material (such as a liquid solution or a crystalline solution) occurring naturally as a product of inorganic processes. Most minerals are solids. Water and mercury are examples of liquid minerals, and air and helium (from rocks or helium wells) are examples ofgaseous minerals. Amalgam (mercury containing dissolved silver and gold) is an example of a solution occurring as a mineral. Rocks are simple minerals (limestone consists of the mineral calcite, which is calcium carbonate) or mixtures of minerals (granite is such a mixture).


A substance is usually defined by chemists as a homogeneous species of matter with reasonably definite chemical composition.

By this definition, pure salt, pure sugar, pure iron, pure copper, pure sulfur, pure water, pure oxygen, and pure hydrogen are representative substances. On the other hand, a solution of sugar in water is not a substance; it is, to be sure, homogeneous, but it does not satisfy the second part of the above definition, inasmuch as its composition is not definite but is widely variable, being determined by the amount of sugar that happens to have been dissolved in a given amount of water. Similarly, the gold of a gold ring or watchcase is not a pure substance, even though it is apparently homogeneous. It is an alloy of gold with other metals, and it usually consists of a crystalline solution of copper in gold. The word alloy is used to refer to a metallic material containing two or more elements: the intermetallic compounds are substances, but most alloys are crystalline solutions or mixtures.


Excerpted from General Chemistry by Linus Pauling. Copyright © 1970 Linus Pauling. Excerpted by permission of Dover Publications, Inc..
All rights reserved. No part of this excerpt may be reproduced or reprinted without permission in writing from the publisher.
Excerpts are provided by Dial-A-Book Inc. solely for the personal use of visitors to this web site.

Table of Contents

1 The Nature and Properties of Matter
1-1 Matter and Chemistry
1-2 Mass and Energy
1-3 The International System of Units
1-4 Temperature
1-5 Kinds of Matter
1-6 The Physical Properties of Substances
1-7 The Chemical Properties of Substances
1-8 The Scientific Method
2 The Atomic and Molecular Structure of Matter
2-1 "Hypotheses, Theories, and Laws"
2-2 The Atomic Theory
2-3 Modern Methods of Studying Atoms and Molecules
2-4 The Arrangement of Atoms in a Crystal
2-5 The Description of a Crystal Structure
2-6 Crystal Symmetry; the Crystal Systems
2-7 The Molecular Structure of Matter
3 "The Electron, the Nuclei of Atoms, and the Photon"
3-1 The Nature of Electricity
3-2 The Discovery of the Electron
3-3 The Discovery of of X-rays and Radioactivity
3-4 The Nuclei of Atoms
3-5 The Birth of the Quantum Theory
3-6 The Photoelectric Effect and the Photon
3-7 The Diffraction of X-rays by Crystals
3-8 Electron Wave Character and Electron Spin
3-9 What Is Light? What Is an Electron?
3-10 The Uncertainty Principle
4 Elements and Compounds. Atomic and Molecular Masses
4-1 The Chemical Elements
4-2 The Neutron. The Structure of Nuclei
4-3 Chemical Reactions
4-4 Nuclidic Masses and Atomic Weights
4-5 Avogadro's Number. The Mole
4-6 Examples of Weight-relation Calculations
4-7 Determination of Atomic Weights by Chemical Method
4-8 Determination of Atomic Weights by Use of the Mass Spectrograph
4-9 Determination of Nuclidic Masses by Nuclear Reactions
4-10 The Discovery of the Correct Atomic Weights. Isomorphism
5 Atomic Structure and the Periodic Table of the Elements
5-1 The Bohr Theory of the Hydrogen Atom
5-2 Excitation and Ionization Energies
5-3 The Wave-mechanical Description of Atoms
5-4 The Periodic Table of the Elements
5-5 Electron Energy as the Basis of the Periodic Table
5-6 The History of the Periodic Table
6 The Chemical Bond
6-1 The Nature of Covalence
6-2 The Structure of Covalent Compounds
6-3 The Direction of Valence Bonds in Space
6-4 Tetrahedral Bond Orbitals
6-5 Bond Orbitals with Large p Character
6-6 Molecules and Crystals of the Nonmetallic Elements
6-7 Resonance
6-8 Ionic Valence
6-9 The Partial Ionic Character of Covalent Bonds
6-10 The Electronegativity Scale of the Elements
6-11 Heats of Formation and Relative Electronegativity of Atoms
6-12 The Electroneutrality Principle
6-13 The Sizes of Atoms and Molecules.
Covalent Radii and van der Waals Radii
6-14 Oxidation Numbers of Atoms
7 The Nonmetallic Elements and Some Their Compounds
7-1 The Elementary Substances
7-2 Hydrides of Nonmetals. Hydrocarbons
7-3 Hydrocarbons Containing Double Bonds and Triple Bonds
7-4 Aromatic Hydrocarbons. Benzene
7-5 Amnonia and Its Compounds
7-6 Other Normal-valence Componds of the Nonmetals
7-7 Some Transargononic Single-bonded Compounds
7-8 The Argonons
8 Oxygen Compounds of Nonmetallic Elements
8-1 The Oxycompounds of the Halogens
8-2 "Oxycompounds of Sulfur, Selenium, and Tellurim"
8-3 "Oxycompounds of Phosphorus, Arsenic, Antimony, and Bismuth"
8-4 Oxycompounds of Nitrogen
8-5 Oxycompounds of Carbon
8-6 Molecules containing Bivalent Carbon. Free Radicals
8-7 Unstable and Highly Reactive Molecules
9 Gases: Quantum Mechanics and Statistical Mechanics
9-1 The Perfect-gas Equation
9-2 Quantum Mechanics of a Monatomic Gas
9-3 The Wave Equation
9-4 The Kinetic Theory of Gases
9-5 The Distribution Law for Molecular Velocities
9-6 The Boltzmann Distribution Law
9-7 Deviations of Real Gases from Ideal Behavior
10 Chemical Thermodynamics
10-1 Heat and Work. Energy and Enthalpy
10-2 The First Law of Thermodynamics
10-3 "Heat Capacity. Heats of Fusion, Vaporization, and Transition"
10-4 Entropy. The Probable State of an Isolated System
10-5 The Absolute Entropy of a Perfect Gas
10-6 Reversible and Irreversible Changes in State
10-7 The Efficiency of a Heat Engine
10-8 Change in Entropy of Any System with Temperature
10-9 The Third Law of Thermodynamics
10-10 The Heat Capacity of Diatomic Gases
10-11 Quantum States of the Rigid Rotator
10-12 The Rotational Entropy of Diatomic Gases
10-13 Quantum States of the Harmonic Oscillator
10-14 Vibrational States of Diatomic Molecules
10-15 "Energy, Heat Capacity, and Entropy of a Harmonic Oscillator"
10-16 The Quantum Theory of Low-temperature Heat Capacity of Crystals
11 Chemical Equilibrium
11-1 The Thermodynamic Conditon for Chemical Equilibrium
11-2 The Vapor Pressure of a Liquid or Crystal
11-3 "Entropy of Transition, Fusion, and Vaporization"
11-4 Van der Waals Forces. Melting Points and Boiling Points
11-5 Chemical Equilibrium in Gases
11-6 Change of Equilibrium with Temperature
11-7 Equilibrium in Heterogeneous Systems
11-8 Le Chatelier's Principle
11-9 The Phase Rule-a Method of Classifying All Systems in Equilibrium
11-10 The Conditions under Which a Reaction Proceeds to Completion
12 Water
12-1 The Composition of Water
12-2 The Water Molecule
12-3 The Properties of Water
12-4 The Hydrogen Bond-the Cause of the Unusual Properties of Water
12-5 The Entropy of Ice
12-6 The Importance of Water as an Electrolytic Solvent
12-7 Heavy Water
12-8 Deviation of Water and Some Other Liquids from Hildebrand's Rule
12-9 The Dense Forms of Ice
12-10 The Phase Diagram of Water
13 The Properties of Solutions
13-1 Types of Solutions. Nomenclature
13-2 Solubility
13-3 The Dependence of Solubility on the Nature of Solute and Solvent
13-4 Solubility of Salts and Hydroxides
13-5 The Solubility-Product Principle
13-6 The Solubility of Gases in Liquids: Henry's Law
13-7 The Freezing Point and Boiling Point of Solution
13-8 The Vapor Pressure of Solutions: Raoult's Law
13-9 The Osmotic Pressure of Solutions
13-10 The Escaping Tendency and the Chemical Potential
13-11 The Properties of Ionic Solutions
13-12 Colloidal Solutions
14 Acids and Bases
14-1 Hydronium-ion (Hydrogen-ion) Concentration
14-2 The Equilibrium between Hydrogen Ion and Hydroxide Ion in Aqueous Solution
14-3 Indicators
14-4 Equivalent Weights of Acids and Bases
14-5 Week Acids and Bases
14-6 The Titration of Weak Acids and Bases
14-7 Buffered Solutions
14-8 The Strengths of the Oxygen Acids
14-9 The Solution of Carbonates in Acid; Hard Water
14-10 The Precipitation of Sulfides
14-11 Nonaqueous Amphiprotic Solvents
15 Oxidation-Reduction Reactions. Electrolysis
15-1 The Electrolytic Decomposition of Molten Salts
15-2 The Electrolysis of and Aqueous Salt Solution
15-3 Oxidation-Reduction Reactions
15-4 Quantitative Relations in Electrolysis
15-5 The Electromotive-force Series of the Elements
15-6 Equilibrium Constants for Oxidation-Reduction Couples
15-7 The Dependence of the Elctomotive Force of Cells on Concentration
15-8 Primary Cells and Storage Cells
15-9 Electrolytic Production of Elements
15-10 The Reduction of Ores. Metallurgy
16 The Rate of Chemical Reactions
16-1 Factors Influencing the Rate of Reactions
16-2 The Rate of a First-order Reaction at Constant Temperature
16-3 Reactions of Higher Order
16-4 Mechanism of Reactions. Dependence of Reaction Rate on Temperature
16-5 Catalysis
16-6 Kinetics of Enzyme Reactions
16-7 Chain Reactions
17 The Nature of Metals and Alloys
17-1 The Metallic Elements
17-2 The Structure of Metals
17-3 The Nature of the Transition Metals
17-4 The Metallic State
17-5 Metallic Valence
17-6 The Free-electron Theory of Metals
17-7 The Nature of Alloys
17-8 Experimental Methods of Studying Alloys
17-9 Interstitial Solid Solutions and Substitutional Solid Solutions
17-10 Physical Metallurgy
18 "Lithium, Beryllium, Boron, and Silicon and Their Congeners"
18-1 "The Electronic Structures of Lithium, Beryllium, Boron, and Silicon and Their Congeners"
18-2 "Radius Ratio, Ligancy, and the Properties of Substances"
18-3 The Alkali Metals and Their Compounds
18-4 The Alkaline-earth Metals and Their Compounds
18-5 Boron
18-6 "The Boranes, Electron-deficient Substances"
18-7 Aluminum and Its Congeners
18-8 Silicon and Its Simpler Compounds
18-9 Silicon Dioxide
18-10 Sodium Silicate and other Silicates
18-11 The Silicate Minerals
18-12 Glass
18-13 Cement
18-14 The Silicones
18-15 Germanium
18-16 Tin
18-17 Lead
19 Inorganic Complexes and the Chemistry of the Transiton Metals
19-1 The Nature of Inorganic Complexes
19-2 "Tetrahedral, Octahedral, and Square Bond Orbitals"
19-3 Ammonia Complexes
19-4 Cyanide Complexes
19-5 Complex Halides and Other Complex Ions
19-6 Hydroxide Complexes
19-7 Sulfide Complexes
19-8 The Quantitive Treatment of Complex Formation
19-9 Polydentate Complexing Agents
19-10 The Structure and Stability of Carbonyls and Other Covalent Complexes of the Transition Metals
19-11 Polynuclear Complexes
20 "Iron, Cobalt, Nickel, and the Platinum Metals"
20-1 "The Electronic Structures and Oxidation States of Iron, Cobalt, Nickel, and the Platinum Metals"
20-2 Iron
20-3 Steel
20-4 Compounds of Iron
20-5 Cobalt
20-6 Nickel
20-7 The Platinum Metals
21 "Copper, Zinc, and Gallium and Their Congeners"
21-1 "The Electronic Structures and Oxidation States of Copper, Silver, and Gold"
21-2 "The Properties of Copper, Silver, and Gold"
21-3 The Compounds of Copper
21-4 The Compounds of Silver
21-5 Photochemistry and Photography
21-6 The Compounds of Gold
21-7 Color and Mixed Oxidation States
21-8 "The Properties and Uses of Zinc, Cadmium, and Mercury"
21-9 Compounds of Zinc and Cadmium
21-10 Compounds of Mercury
21-11 "Gallium, Indium, and Thallium"
22 "Titanium, Vanadium, Chromium, and Manganese and Their Congeners"
22-1 "The Electronic Structures of Titanium, Vanadium, Chromium, and Manganese and Their Congeners"
22-2 "Titanium, Zirconium, Hafnium, and Thorium"
22-3 "Vanadium, Niobium, Tantalum, and Protactinium"
22-4 Superconductivity
22-5 Chromium
22-6 The Congeners of Chromium
22-7 Managanese
22-8 Acid-forming and Base-forming Oxides and Hydroxides
22-9 The Congeners of Manganese
23 Organic Chemistry
23-1 The Nature and Extent of Organic Chemistry
23-2 Petroleum and the Hydrocarbons
23-3 Alcohols and Phenols
23-4 Aldehydes and Ketones
23-5 The Organic Acids and Their Esters
23-6 Amines and Other Organic Compounds of Nitrogen
23-7 "Carbohydrates, Sugars, Polysaccharides"
23-8 Fibers and Plastics
24 Biochemistry
24-1 The Nature of Life
24-2 The Structure of Living Organisms
24-3 Amino Acids and Protiens
24-4 Nucleic Acids. The Chemistry of Heredity
24-5 Metabolic Processes. Enzymes and Their Action
24-6 Vitamins
24-7 Hormones
24-8 Chemistry and Medicine
25 The Chemistry of the Fundamental Particles
25-1 The Classification of the Fundemental Particles
25-2 The Discovery of the Fundemental Particles
25-3 The Forces between Nucleons. Strong Interactions
25-4 The Structure of Nucleons
25-5 Leptons and Antileptons
25-6 Mesons and Antimesons
25-7 Baryons and Antibaryons
25-8 The Decay Reactions of the Fundemental Particles
25-9 Strangeness (Xenicity)
25-10 Resonance Particles and Complexes
25-11 The Structure of the Fundamental Particles. Quarks
25-12 "Positronium, Muonium, Mesonic Atoms"
26 Nuclear Chemistry
26-1 Natural Radioactivity
26-2 The Age of the Earth
26-3 Artificial Radioactivity
26-4 The Kinds of Nuclear Reactions
26-5 The Use of Radioactive Elements as Tracers
26-6 Dating Objects by Use of Carbon
26-7 The Properties of Nucleides
26-8 The Shell Model of Nuclear Structure
26-9 The Helion-Triton Model
26-10 Nuclear Fission and Nuclear Fusion
I. Units of Measurement
II. Values of Some Physical and Chemical Constants
III. Symmetry of Molecules and Crystals
IV. X-rays and Crystal Structure
V. Hydrogenlike Orbitals
VI. Russel-Saunders States of Atoms Allowed by the Pauli Exclusion Principle
VII. Hybrid Bond Orbitals
VIII. Bond Energy and Bond-dissociation Energy
IX. The Vapor Pressure of Water
X. An Alternitive Derivation of the Boltzmann Distribution Law
XI. The Boltzmann Dristribution Law in Classical Mechanics
XII. The Entropy of a Perfect Gas
XIII. Electric Polarizabilities and Electric Dipole Moments
XIV. The Magnetic Properties of Substances
XV. Values of Thermodynamic Properties of Some Substances at 25°C and 1 atm
XVI. Selected Readings

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General Chemistry 4.8 out of 5 based on 0 ratings. 4 reviews.
Guest More than 1 year ago
Pauling's writing style is engaging, with just the right mix of technical content, historical treatment, and editorial. The problems included with the text reinforce the subject matter without being too facile. The main drawback is that some of the material is now dated. In particular, some of the physical constants in the text are not longer accurate. Keeping this in mind, this work will serve as a wonderful reference, as well as a great introduction to the subject. As with most Dover publications, an excellent value!!
Anonymous More than 1 year ago
I love this stuff.
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