Concise Physical Chemistry

Donald W. Rogers

• Publisher: Wiley
• Year: 2011
• ISBN:
  9781118102237
• Edition: 1
• Topic: Physical Sciences
• Subtopic: Physical & Theoretical Chemistry

This book is a physical chemistry textbookthat presents the essentials of physical chemistry as a logical sequence from its most modest beginning to contemporary research topics. Many books currently on the market focus on the problem sets with a cursory treatment of the conceptual background and theoretical material, whereas this book is concerned only with the conceptual development of the subject.Comprised of 19 chapters, the book will address ideal gas laws, real gases, the thermodynamics of simple systems, thermochemistry, entropy and the second law, the Gibbs free energy, equilibrium, statistical approaches to thermodynamics, the phase rule, chemical kinetics, liquids and solids, solution chemistry, conductivity, electrochemical cells, atomic theory, wave mechanics of simple systems, molecular orbital theory, experimental determination of molecular structure, and photochemistry and the theory of chemical kinetics.

CHAPTER 1

IDEAL GAS LAWS

In the seventeenth and eighteenth centuries, thoughtful people, influenced by the success of early scientists like Galileo and Newton in the fields of mechanics and astronomy, began to look more carefully for quantitative connections among the phenomena around them. Among these people were the chemist Robert Boyle and the famous French balloonist Jacques Alexandre César Charles.

1.1 EMPIRICAL GAS LAWS

Many physical chemistry textbooks begin, quite properly, with a statement of Boyle’s and Charles’s laws of ideal gases:

and

The constants k1 and k2 can be approximated simply by averaging a series of experimental measurements, first of pV at constant temperature T for the Boyle equation, then of V/T at constant pressure p for Charles’s law. All this can be done using simple manometers and thermometers.

1.1.1 The Combined Gas Law

These two laws can be combined to give a new constant

Subsequently, it was found that if the quantity of gas taken is the number of grams equal to the atomic or molecular weight of the gas, the constant k3, now written R under the new stipulations, is given by

For the number of moles of a gas, n, we have

The constant R is called the universal gas constant.

1.1.2 Units

The pressure of a confined gas is the sum of the force exerted by all of the gas molecules as they impact with the container walls of area A in unit time:

The summed force f is given in units of newtons (N), and the area is in square meters (m2). The N m− 2 is also called the pascal (Pa). The pascal is about five or six orders of magnitude smaller than pressures encountered in normal laboratory practice, so the convenient unit 1 bar ≡ 105 Pa was defined.

The logical unit of volume in the MKS (meter, kilogram, second) system is the m3, but this also is not commensurate with routine laboratory practice where the liter is used. One thousand liters equals 1 m3, so the MKS name for this cubic measure is the cubic decimeter—that is, one-tenth of a meter cubed (1 dm3). Because there are 1000 cubic decimeters in a cubic meter and 1000 liters in a cubic meter, it is evident that 1 L = 1 dm3.

The unit of temperature is the kelvin (K), and the unit of weight is the kilogram (kg). Formally, there is a difference between weight and mass, which we shall ignore for the most part. Chemists are fond of expressing the amount of a pure substance in terms of the number of moles n (a pure, unitless number), which is the mass in kg divided by an experimentally determined unit molar mass M, also in kg:1

If the pressure is expressed as N m− 2 and volume is in m3, then pV has the unit N m, which is a unit of energy called the joule (J). From this, the expression

gives the unit of R as J K− 1 mol− 1. Experiment revealed that

which also defines the atmosphere, an older unit of pressure that still pervades the literature.

1.2 THE MOLE

The concept of the mole (gram molecular weight in early literature) arises from the deduction by Avogadro in 1811 that equal volumes of gas at the same pressure and temperature contain the same number of particles. This somewhat intuitive conclusion was drawn from a picture of the gaseous state as being characterized by repulsive forces between gaseous particles whereby doubling, tripling, and so on, the weight of the sample taken will double, triple, and so on, its number of particles, hence its volume. It was also known at the time that electrolysis of water produced two volumes of hydrogen for every volume of oxygen, so Avogadro deduced the formula H2 O for water on the basis of his hypothesis of equal volume for equal numbers of particles in the gaseous state.

By Avogadro’s time, it was also known that the number of grams of oxygen obtained by electrolysis of water is 8 times the number of grams of hydrogen. By his 2-for-1 hypothesis, Avogadro reasoned that the less numerous oxygen a...

Table of contents

• Cover
• Half Title page
• Title page
• Copyright page
• Foreword
• Preface
• Chapter 1: Ideal Gas Laws
• Chapter 2: Real Gases: Empirical Equations
• Chapter 3: The Thermodynamics of Simple Systems
• Chapter 4: Thermochemistry
• Chapter 5: Entropy and the Second Law
• Chapter 6: The Gibbs Free Energy
• Chapter 7: Equilibrium
• Chapter 8: A Statistical Approach to Thermodynamics
• Chapter 9: The Phase Rule
• Chapter 10: Chemical Kinetics
• Chapter 11: Liquids and Solids
• Chapter 12: Solution Chemistry
• Chapter 13: Coulometry and Conductivity
• Chapter 14: Electrochemical Cells
• Chapter 15: Early Quantum Theory: A Summary
• Chapter 16: Wave Mechanics of Simple Systems
• Chapter 17: The Variational Method: Atoms
• Chapter 18: Experimental Determination of Molecular Structure
• Chapter 19: Classical Molecular Modeling
• Chapter 20: Quantum Molecular Modeling
• Chapter 21: Photochemistry and the Theory of Chemical Reactions
• References
• Answers to Selected Odd-Numbered Problems
• Index