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SLOW STEPS BALANCING CHEMICAL EQUATIONS WITH LEAST COMMON MULTIPLES
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Chemistry: TheCentral Science - 8th Edition

Index of Chapter Contents

 

1- Introduction: Matter and Measurement

1.1 The Study of Chemistry

In this introductory section you will learn a few basic terms and get a very brief preview of what lies ahead in this chapter.

1.2 Classification of Matter

This section introduces the states of matter: gas, liquid, and solid. It also provides information about the nature of each phase. Most matter exists in mixtures that can be separated into simpler substances. You will learn how to classify a sample of matter as a mixture, a compound, or an element.

1.3 Properties of Matter

Different samples of matter have different physical and chemical properties. This section explores physical and chemical properties of some elements and the physical and chemical processes that can be used to determine those properties.

1.4 Units of Measurement

The scientific community requires a system of units that can be used worldwide to report experimental results. These preferred units are called SI units; the SI units and the decimal prefixes used to denote magnitude are introduced in this section.

1.5 Uncertainty in Measurement

Regardless of how carefully measurements are made, numbers obtained by measurement are always inexact. In this section you will learn the terms "accuracy" and "precision," and the difference between them. You will also learn how to report measured and calculated numbers using significant figures to indicate the exactness of a measurement.

1.6 Dimensional Analysis

Essential to success in problem solving, dimensional analysis is introduced in this section. You will see the importance of writing the units associated with numbers when solving problems, and you will solve some problems using conversion factors.


2 - Atoms, Molecules, and Ions

2.1 The Atomic Theory of Matter

The first meaningful atomic theory was put forth early in the nineteenth century. In this section you will see how this atomic theory explained some well-known scientific laws and enabled Dalton to predict the law of multiple proportions. You will see an animation of a combustion reaction that illustrates the law of multiple proportions.

2.2 The Discovery of Atomic Structure

Here you will learn about some of the historic experiments that led to our knowledge of atomic structure. Included are animations of the Millikan oil-drop experiment, Rutherford's gold foil experiment, and experiments with radioactivity.

2.3 The Modern View of Atomic Structure

Only three subatomic particles have an effect on chemical behavior: protons, neutrons, and electrons. This section gives the charges and masses of these three subatomic particles, and introduces isotopes-atoms of a single element having different numbers of neutrons.

2.4 The Periodic Table

Elements are arranged in the periodic table in order of increasing atomic number. This section introduces the periodic table and the ways in which elements can be categorized according to their properties.

2.5 Molecules and Molecular Compounds

The atom is the smallest unit of matter, but most matter does not exist as isolated atoms. Many substances exist as molecules that consist of two or more atoms. Various ways of representing molecular substances are presented in this section, including molecular and empirical formulas.

2.6 Ions and Ionic Compounds

Many substances exist as a combination of ions, charged atoms, or collections of atoms. Ionic compounds do not consist of molecules, and so do not have molecular formulas. Ionic substances are represented with empirical formulas. The importance of knowing the names, formulas, and charges of common polyatomic ions is stressed.

2.7 Naming Inorganic Compounds

The basic rules of nomenclature are introduced for naming ionic compounds and molecular compounds, including acids. An interactive tutorial will allow you to practice deducing formulas and names for ionic compounds.


3 - Stoichiometry: Calculations with Chemical Formulas and Equations

3.1 Chemical Equations

A chemical equation is a written representation of a chemical reaction. In this section you will learn how to interpret and balance chemical equations. You will also have the opportunity to develop your skills balancing equations using an interactive tutorial.

3.2 Patterns of Chemical Reactivity

At first glance the many different reactions you study in chemistry may seem unrelated. However, a closer look reveals distinct patterns of chemical behavior among certain groups of elements. You will learn some of these patterns of chemical reactivity, and you will view demonstrations of several examples.

3.3 Atomic and Molecular Weights

Because atoms and molecules have masses too small to express conveniently in conventional units, we use the atomic mass scale to express their masses in units called atomic mass units (amu). The amu is an extremely small mass; 1 amu is equal to 1.66054 10-24 g. You will learn to compute molecular weight and percentage composition using the atomic mass scale.

3.4 The Mole

Because atoms and molecules are so small, the number of such particles in a macroscopic sample is extremely large. For example, an 18-ml sample of water contains 6.02 1023 water molecules; an unimaginably large number. However, since the quantities of substances used in chemical reactions are typically at least this big, it is necessary to have a way to keep track of the enormous numbers of molecules involved. In chemistry the unit for dealing with atoms, ions, and molecules is the mole, abbreviated mol. A mole is simply a number of items - in this case molecules. One mol of water contains 6.02 1023 water molecules. Another example: One mol of 12C atoms is 6.02 1023 atoms.

3.5 Empirical Formulas from Analyses

The empirical formula for a substance tells us the relative number of atoms of each element it contains. In this section you will learn how to determine the empirical formula of a compound from its percentage composition, which is often determined by combustion analysis.

3.6 Quantitative Information from Balanced Equations

In this section you will learn to manipulate a balanced chemical equation in order to make quantitative predictions such as how much product will be formed from given quantities of reactants. Proficiency at converting from grams to moles and from moles to grams is essential to your understanding of this section.

3.7 Limiting Reactants

The amount of product obtained from a chemical reaction is nearly always limited by the available amount of one of the reactants. The limiting reactant is the one that is completely consumed in the process. After one of the reactants has been depleted, the reaction stops and no more product can be formed. In this section you will learn how to determine which reactant is the limiting reactant; you will also learn how to measure the efficiency of a reaction by calculating the percent yield.


4 - Aqueous Reactions and Solution Stoichiometry

4.1 General Properties of Aqueous Solutions

Solutions are simply homogeneous mixtures. All solutions consist of a solvent and one or more solutes. Solutes may be electrolytes or nonelectrolytes. In this section you will see an animation of salt and sugar solutions used to complete an electrical circuit. You will also see an animation that illustrates the difference between strong and weak electrolytes.

4.2 Precipitation Reactions

Combinations of solutions can sometimes produce a solid product called a precipitate. In this section you will learn how to write complete ionic and net ionic reactions, how to predict the products of a chemical reaction, and how to tell whether a solid product will form.

4.3 Acid-Base Reactions

In this section you will see two more animations that introduce aqueous acids and bases. Reactions between acids and bases are also introduced.

4.4 Oxidation-Reduction Reactions

Oxidation-reduction reactions involve the transfer of electrons from one species to another. Oxidation numbers are introduced. You will see an animation and a demonstration of oxidation-reduction reactions, and you will learn to predict what reactions will occur, based on the activity series

4.5 Concentrations of Solutions

Molarity is the most commonly used expression of solution concentration. Using an interactive activity, you will learn how to calculate it, and you will see animations of solutions being prepared both from solid solute and from concentrated stock solutions.

4.6 Solution Stoichiometry and Chemical Analysis

Solution stoichiometry facilitates analytical techniques such as titration. In this section you will see an animation of the titration process and learn how to use data from a titration to calculate the concentration of an unknown acid. You will practice your titration technique with an interactive simulation.


5 - Thermochemistry

5.1 The Nature of Energy

Thermochemistry is the study of the energy changes associated with chemical and physical processes. In this section you will learn about two types of energy that matter can possess: potential energy and kinetic energy. In addition, you will see the distinction between a system and its surroundings in a thermochemical experiment.

5.2 The First Law of Thermodynamics

The first law of thermodynamics is that energy is conserved. You will see an example of energy being converted from potential to kinetic, and vice versa. The terms exothermic and endothermic are introduced. You will have the chance to view a demonstration of an exothermic reaction, and you will be able to explore an endothermic process with an interactive simulation. You will learn that internal energy (E) of a system cannot be measured, but that changes in internal energy ( E) can be measured and that the change in a system's internal energy is equal to the sum of heat transferred to the system and work done on the system.

5.3 Enthalpy

The term enthalpy is introduced in this section. Related to energy, the enthalpy (H) of a system cannot be measured, but a change in enthalpy ( H) can be measured. A system's change in enthalpy is defined as the amount of heat transferred between system and surroundings during a process that occurs at constant pressure. Since many reactions and processes of interest are carried out at constant pressure, the concept of enthalpy is an important one.

5.4 Enthalpies of Reaction

In this section you will see that there is an enthalpy change that accompanies a reaction. It is called the enthalpy of reaction, or the heat of reaction, and can be written as Hrxn. Enthalpy of reaction values are associated with many different reactions and processes. You will use an interactive simulation to learn specifically about enthalpies of dissolution.

5.5 Calorimetry

Thermochemical experiments are carried out using calorimetry, which allows the measuring of heat flow. In this section you will learn the terms and concepts essential to understanding calorimetry, and you will perform simulated calorimetry experiments under conditions of constant volume.

5.6 Hess's Law

Hess's law allows us to determine change in enthalpy for a process that is made up of several steps—provided we know the enthalpy change associated with each individual step. You will see an example of Hess's law being used to compute the enthalpy change of a reaction.

5.7 Enthalpies of Formation

In this section you will be introduced to the concept of enthalpy of formation. The enthalpy of formation is strictly defined as the enthalpy change for a process by which one mole of a substance is formed from its constituent elements, each in their standard states. The standard state for thermodynamics (thermochemistry) is defined as 298 K and 1 atm pressure.

5.8 Foods and Fuels

The concept of enthalpy of reaction is applied to the metabolism of foods and the combustion of fuels. The term "fuel value" is introduced, and you will calculate the calorie content of some common foods.


6 - Electronic Structure of Atoms

6.1 The Wave Nature of Light

Much of what we know about matter comes from experiments involving the interaction of light with matter. Basic properties of light and waves are introduced in this section, and the relationships among wavelength, frequency, and speed of light are presented.

6.2 Quantized Energy and Photons

Early in the twentieth century, energy was thought to be "continuous." Some revolutionary theories, however, changed forever the way we view energy. Energy is now known to be "quantized"—released or absorbed by atoms in "chunks" of some minimum size—and to exhibit certain particle-like behavior. You will see an animation of the photoelectric effect, and youll use an activity designed to illustrate the relationship between the wavelength of light and the energy associated with it.

6.3 Bohrs Model of the Hydrogen Atom

The revolutionary theories involving energy led to other theories involving the nature of the electron and structure of the atom. Bohrs model of the hydrogen atom is presented here. You will see flame tests for elements, and you will learn to relate the colors given off by an element to its electronic structure.

6.4 The Wave Behavior of Matter

In this section you will learn that just as energy can be shown to behave like a stream of particles, tiny particles such as electrons can be made to diffract. Diffraction is a property of waves.

6.5 Quantum Mechanics and Atomic Orbitals

This section introduces the quantum mechanical model of the atom. You will learn about quantum numbers, shells, subshells, and orbitals within the hydrogen atom.

6.6 Representations of Orbitals

Building on the quantum mechanical model, models of the hydrogen atom orbitals are constructed to aid in visualization of atomic structure.

6.7 Orbitals in Many-Electron Atoms

Here, Section 6.5 is extended to apply to atoms and ions with more than one electron. The relative energies of orbitals change and now depend on the value of the second quantum number in addition to the principal quantum number.

6.8 Electron Configurations

Using the knowledge of relative orbital energies, we are able to predict the electron configuration for any element in the periodic table. Hunds rule and the Pauli exclusion principle are presented, and you will be able to practice predicting electron configurations and writing the appropriate spectroscopic notation.

6.9 Electron Configurations and the Periodic Table

Finally, the prediction of electron configurations is simplified by dividing the periodic table into blocks associated with a particular subshell. The s-block elements have valence electrons in an s orbital. The p-block elements have valence electrons in a p orbital, and so on.


7 - Periodic Properties of the Elements

7.1 Development of the Periodic Table

The periodic table is a very important tool used by chemists to organize information about the elements. A very brief history of the modern periodic table is presented along with a picture of Mendeleev's early periodic table. You will also see a modern interactive periodic table.

7.2 Electron Shells and the Sizes of Atoms

The size of an atom is defined, and the sizes of atoms are compared to point out one trend in the elements. You will see an animation that illustrates trends in atomic size, both across a period (row) and down a group (column) in the periodic table.

7.3 Ionization Energy

The reactivity of an element depends in part on its ionization energy. This section contains two animations: one of the ionization process and one that illustrates the trends in ionization energy across and down the periodic table. You will also use a module to graph ionization energies versus atomic number.

7.4 Electron Affinities

Another important factor in reactivity, electron affinity, is defined. This section also contains two animations: one describing electron affinity and one illustrating the periodic trends in electron affinity.

7.5 Metals, Nonmetals, and Metalloids

In this section the chemical properties of metals and nonmetals are presented. Two animations illustrate the chemical behavior of oxides (metallic and nonmetallic) and general periodic trends, including metallic character.

7.6 Group Trends for the Active Metals

Some of the physical and chemical properties of the alkali metals and alkaline earth metals are presented here. You will also watch a demonstration that shows the difference in reactivity of two alkali metals.

7.7 Group Trends for Selected Nonmetals

Hydrogen, the oxygen group, the halogens, and the noble gases are all briefly presented in this section. You will watch a movie that illustrates some of the physical properties of the halogens.


8 - Basic Concepts of Chemical Bonding

8.1 Chemical Bonds, Lewis Symbols, and the Octet Rule

Three different types of chemical bonds are introduced: ionic, covalent, and metallic. Ionic and covalent bonds are the focus of this chapter. Metallic bonding will be discussed in detail in eChapter 23.5. Here, you will see that atoms tend to lose, gain, or share electrons to attain a noble-gas electron configuration. This observation is summarized in the octet rule. You will learn to draw Lewis symbols to represent atoms and ions and their valence electrons.

8.2 Ionic Bonding

This section focuses on the bonding due to electrostatic attraction between oppositely charged particles. You will see a movie in which the elements sodium and chlorine react exothermically to form sodium chloride. The principle of electrostatic attraction will be reviewed with an activity in which you can vary the charges on, and distance between, two particles to see what effect each has on the magnitude of the force between them. The term lattice energy is introduced as a quantitative measure of the attractive forces between ions in a solid.

8.3 Sizes of Ions

Electron configurations and electrostatic forces are used to explain the differences in size between atoms and their corresponding ions. You will watch a movie of the electron-loss and electron-gain processes that details how each affects size.

8.4 Covalent Bonding

Some atoms share electrons to achieve a noble-gas electron configuration. A chemical bond formed by sharing a pair of electrons is called a covalent bond. You will see a movie of the process by which two hydrogen atoms form a covalent bond. Electrostatic forces again are used to explain the length of the H–H bond.

8.5 Bond Polarity and Electronegativity

This section introduces the concept of electronegativity and illustrates the trends in electronegativity contained in the periodic table. Electronegativity differences give rise to polar bonds, which are bonds in which electron density between atoms is not shared equally. You will learn to calculate a dipole moment—a quantitative measure of the polarity of a bond. An activity will represent bond dipoles as vectors for bonds between atoms that you choose.

8.6 Drawing Lewis Structures

The rules for representing molecules and polyatomic ions with Lewis structures are introduced here. You will determine which bonds in a molecule are single bonds, and which are double or triple bonds. A movie explains formal charges, a method of keeping track of electrons that helps you determine the best Lewis structures. Nomenclature is revisited briefly from the standpoint of using the covalent nomenclature convention for compounds that contain metals.

8.7 Resonance Structures

Often, there are two or more equally correct Lewis structures for the same species. Those that differ only by placement of electrons are called resonance structures. None of the resonance structures truly represents the molecule accurately. The bond lengths and strengths are intermediate between the possible resonance structures.

8.8 Exceptions to the Octet Rule

There are three types of exceptions to the octet rule. All three are discussed in this section and examples are given for each. You will again use formal charges to decide which structures are most feasible.

8.9 Strengths of Covalent Bonds

Bond enthalpy is defined in this section, and tabulated values of average bond enthalpies are used to calculate enthalpies of reaction. The strength of a bond and its length are shown to have a reciprocal relationship.


9 - Molecular Geometry and Bonding Theories

9.1 Molecular Shapes

The size and shape of a molecule of a particular substance play an important part in determining the physical and chemical properties of that substance. Although Lewis structures are quite useful, they do not indicate a molecule's shape. In this section the Lewis structure for methane will be compared with other representations of the methane molecule.

9.2 The VSEPR Model

Based on repulsion between like charges, the VSEPR model helps us account for the arrangements of atoms in polyatomic molecules. The term electron domain is introduced, and you will see an animation that details the arrangements of electron domains about a central atom.

9.3 Polarity of Polyatomic Molecules

Polarity is revisited in the context of polyatomic molecules. A simulation will show how to determine the net dipole moment, if any, of a molecule containing polar bonds.

9.4 Covalent Bonding and Orbital Overlap

The bonding in simple molecules such as H2 and Cl2 is explained by overlap of singly occupied atomic orbitals. This is the introduction to valence-bond theory.

9.5 Hybrid Orbitals

Simple valence-bond theory fails to explain how atoms with no unpaired electrons can form covalent bonds. The use of hybrid orbitals shows us how bonds can form in such molecules as BeF2 and BCl3. An animation illustrates the hybridization of nitrogen orbitals in ammonia.

9.6 Multiple Bonds

Lewis structures sometimes contain double or triple bonds. In this section you will see how the bonds are formed by a combination of hybrid orbitals and unhybridized atomic orbitals.

9.7 Molecular Orbitals

Certain properties of molecules cannot be explained with valence-bond theory or with orbital hybridization. A new theory is presented to explain some of those properties. You will learn the term bond order and how to calculate it as a measure of a bond's stability.

9.8 Second-Row Diatomic Molecules

The molecular orbital energy diagrams for homonuclear diatomic molecules are given. The ordering of orbital energies is explained in terms of orbital interaction. The MO energy-level diagram for oxygen predicts that it should have two unpaired electrons and should be paramagnetic, which it is. This is one of the things that valence-bond theory could not explain. The progression from Lewis structures to VSEPR to valence-bond theory to hybridization and finally to molecular orbital theory is an outstanding example of the scientific method at work.


10 - Gases

10.1 Characteristics of Gases

Gases are one of the three states of matter. A gas is compressible and will assume the shape and volume of its container. This eChapter focuses on the characteristics and behavior of gases.

10.2 Pressure

The physical condition, or state, of a gas is defined by four variables: temperature, pressure, volume, and the number of moles (the amount of the gas). Pressure is introduced here as are units and methods of measuring pressure.

10.3 The Gas Laws

This section presents the relationships between volume and pressure (Boyle's law), volume and temperature (Charles's law), and volume and number of moles (Avogadro's law). An animation illustrates the relationship between volume and pressure at constant temperature, and a simulation lets you explore the relationships described by all three gas laws.

10.4 The Ideal-Gas Equation

Combining the three gas laws yields a general equation for solving gas problems, PV = nRT. Although the equation strictly holds only for ideal gases, it is adequate to solve most gas problems.

10.5 Further Applications of the Ideal-Gas Equation

The ideal-gas equation can be used to solve a variety of problems involving gases. Here, an expression is derived to calculate the density of a gas, and a simulation lets you explore the relationship between molar mass, temperature, pressure, and density. An animation of the process by which an air bag deploys is included, and you will gain some experience interconverting volumes and molar amounts in chemical equations.

10.6 Gas Mixtures and Partial Pressures

Dalton's law of partial pressures states that each component of a gaseous mixture exerts a pressure independent of the other components of the mixture and that the total pressure exerted by the mixture is the sum of the partial pressures for all components. The term mole fraction is introduced as a measure of a gas's concentration in a mixture.

10.7 Kinetic-Molecular Theory

While the gas laws predict how gases behave, the kinetic molecular theory explains why gases behave the way they do. Kinetic molecular theory is used to explain the observations that gave rise to the gas laws. You will view a movie illustrating the relationship between molecular speed and absolute temperature, and a simulation helps you visualize the relationship for a variety of gases.

10.8 Molecular Effusion and Diffusion

Effusion (escape through a small opening) and diffusion (mixing) of gases depend upon molar mass. Graham's law of effusion describes the relative rates of effusion for any two gases. Diffusion, somewhat more complex, is illustrated in a demonstration video.

10.9 Real Gases: Deviations from Ideal Behavior

Although the ideal-gas equation is adequate for most gas problems, there are equations that can more closely predict the behavior of real gases. One of them, the van der Waals equation, is presented.


11- Intermolecular Forces, Liquids, and Solids

11.1 A Molecular Comparison of Liquids and Solids

The three states of matter are revisited, and the differences between them reviewed. Liquids and solids are set apart from gases because they are both condensed phases.

11.2 Intermolecular Forces

The forces that hold particles together and that make condensed phases condensed are collectively called intermolecular forces. The different types of forces are introduced, and examples are given of compounds held together by each type.

11.3 Some Properties of Liquids

Two of the properties of a liquid that depend on the magnitude of the intermolecular forces holding it together are viscosity and surface tension. Both terms are defined, and the relationship between intermolecular forces and the magnitudes of viscosity and surface tension are established.

11.4 Phase Changes

Melting, evaporation, and sublimation, and the reverse processes—freezing, condensation, and deposition—are phase changes. Phase changes are accompanied by energy changes. In this section you will see an animation of the process by which a solid melts and then vaporizes as heat is added.

11.5 Vapor Pressure

In this section you will learn about equilibrium vapor pressure and how it depends on intermolecular forces. An animation will show how vapor pressure depends on temperature, and you will use a simulation to explore this relationship.

11.6 Phase Diagrams

In this section you will learn how to read a phase diagram, which is a graphic way to summarize the conditions under which equilibria exist between the different states of matter—solid, liquid, and gas.

11.7 Structures of Solids

Crystalline solids are the focus of this section. They are very highly ordered in three-dimensional arrays of repeating units called unit cells. Various types of unit cells are introduced, and you will see 3-D simulations of some of the types.

11.8 Bonding in Solids

A crystalline solid is one whose atoms, ions, or molecules are ordered in well-defined arrangements. Four different types of crystalline solids are introduced along with characteristic properties and examples of each. You will see 3-D models of ice (a molecular solid) and two allotropes of carbon: diamond, and graphite (network-covalent solids).


12 - Modern Materials

12.1 Liquid Crystals

Liquid crystals are introduced and several different types are presented. Nematic, smectic A, smectic C, and cholesteric liquid crystals are discussed.

12.2 Polymers

Polymers and two types of polymerization are introduced: addition polymerization and condensation polymerization. You will see a demonstration of the formation of a polymer called Nylon 610. Properties of polymers and the techniques used to modify those properties are also discussed.

12.3 Biomaterials

The physical and chemical properties and requirements of materials used in biomedical applications are presented in this section.

12.4 Ceramics

Properties and processing of ceramics are presented in this section. You will see two 3-D models of ceramic materials, one a carbide, the other a superconductor.

12.5 Thin Films

This section introduces the purpose and manufacture of thin films. Three different methods for producing thin films are presented.


13- Properties of Solutions

13.1 The Solution Process

The solution process is divided into three steps: Two are endothermic, and one is exothermic. The relative magnitudes of these three enthalpy changes determine the overall exo- or endothermicity of the solution process. You will see a movie of the dissolution of sodium chloride in water and use a simulation to explore the enthalpies of dissolution for a number of compounds.

13.2 Saturated Solutions and Solubility

The terms saturated, supersaturated, and unsaturated are introduced in the context of solutions. Solubility is defined as the amount of solute needed to form a saturated solution.

13.3 Factors Affecting Solubility

In this section you will learn the effect on solubility of the various types of intermolecular interactions and the effect of changing temperature and pressure. An animation will illustrate the relationship between the partial pressure of a gas and its solubility in a liquid.

13.4 Ways of Expressing Concentration

This section presents six ways to express the concentration of a solution: mass percentage, parts per million, parts per billion, mole fraction, molarity, and molality. Several of these will be important for use later in this eChapter.

13.5 Colligative Properties

Certain properties of solutions depend on the concentration of dissolved particles–not on the nature of the dissolved particles. These are the colligative properties: lowering the vapor pressure, boiling point elevation, freezing point depression, and osmotic pressure. You will learn how to calculate the expected values of these, and also how to use experimentally determined colligative property values to determine molar mass of an unknown. A simulation shows the effect various solutes have on the boiling point and freezing point of a solution.

13.6 Colloids

Colloidal suspensions are at the dividing line between solutions and heterogeneous mixtures. In this section you will learn what a colloid is and how it can be identified as a colloid; and you'll see some examples of various types of colloids.


14 - Chemical Kinetics

14.1 Reaction Rates

The speed at which a chemical reaction occurs is called the reaction rate. Several things have the capacity to affect the rate of a reaction, including reactant concentrations, temperature, presence of a catalyst, and surface area of reactants or catalysts.

14.2 The Dependence of Rate on Concentration

In this section the terms  rate constant and  rate law are introduced. You will learn how the units of the rate constant depend on the overall  order of a reaction. You will also see the dependence of reaction rate on concentration of reactants, and a simulation in this section will explore this relationship.

14.3 The Change of Concentration with Time

Chemical reactions typically slow down as they progress. In this section you will encounter equations that show how the concentration of a reactant varies as a function of time. The term  half-life is introduced, and you will learn to calculate the half-life of a process. The terms  first order, and  second order are presented, and equations are derived for solving problems of each type. You will use a simulation again in this section to explore the time-concentration connection..

14.4 Temperature and Rate

Reactions occur faster at higher temperatures. In this section the quantitative relationship between rate and temperature is presented. You will learn the term  activation energy, and a simulation will illustrate the relationships that are derived.

14.5 Reaction Mechanisms

Some reactions occur in a single step, but many occur in a  series of steps called  elementary processes. In this section you will watch an animation of a  bimolecular process, and you will learn how to tell whether a proposed reaction mechanism is plausible.

14.6 Catalysis

A catalyst increases the rate of a reaction by providing a different mechanism with a lower activation energy. You will learn about homogeneous and heterogeneous catalysis, and you'll see an animation of a catalyzed reaction on the atomic level.


15 - Chemical Equilibrium

15.1 The Concept of Equilibrium

When a reaction takes place, both the forward process (the reaction as we have written it) and the reverse reaction occur. Equilibrium is the point at which both processes are occurring at the same rate, with no net change of reactant or product concentration. Equilibrium can be achieved starting with only reactants, only products, or some of each. You will use a simulation in this section to plot concentrations of reactants and products as equilibrium is achieved.

15.2 The Equilibrium Constant

There is a constant relationship between equilibrium concentrations of reactants and products at a given temperature. You will learn to write equilibrium expressions for chemical reactions, and a simulation will illustrate the significance of an equilibrium constant's magnitude.

15.3 Heterogeneous Equilibria

Some equilibria involve substances all in the same phase. Many others involve substances in different phases. The concentrations of liquids and solids do not appear in equilibrium expressions for a reaction, and you'll see some examples of heterogeneous equilibria and their equilibrium expressions.

15.4 Calculating Equilibrium Constants

Often, only one of the species present in an equilibrium mixture can be measured. We can generally use the stoichiometry of the reaction to deduce the concentrations of the other species in the chemical equation. We can also calculate the equilibrium constant for such a reaction.

15.5 Applications of Equilibrium Constants

Equilibrium constants can be used to predict the equilibrium concentrations of reactants and products. You will learn how to use K values to calculate the composition of an equilibrium mixture for several different situations.

15.6 Le Châtelier's Principle

An important principle throughout chemistry, Le Châtelier's principle is introduced in this section. Two animations illustrate the effects of certain types of stress that can be applied to a system at equilibrium.


16 - Acid-Base Equilibria

16.1. Acids and Bases: A Brief Review

Acid and base behavior in an aqueous medium is briefly reviewed.

16.2 Brønsted-Lowry Acids and Bases

A new perspective on acids and bases is introduced. The Brønsted-Lowry theory allows us to classify reactions as acid-base if the reactions occur in other than aqueous media. The terms hydrogen ion and hydronium ion are discussed and will be used interchangeably with the symbols H+ and H3O+. The terms conjugate acid and conjugate base are introduced.

16.3 The Autoionization of Water

To a very small extent, pure water ionizes to form hydroxide and hydronium ions. The equilibrium constant for this process is introduced, and the relationship between hydronium and hydroxide ion concentrations in aqueous solutions is presented.

16.4 The pH Scale

Hydronium ion concentration (acidity) of a solution is generally stated as a pH value. The pH scale is introduced in this section. A simulation is used to measure the pH of various acid and base solutions, and a movie shows how common substances can be used as acid-base indicators to estimate the pH of a solution.

16.5 Strong Acids and Bases

In this section you will view two animations. One shows the ionization of three acids, two strong and one weak. The other shows the dissociation of two strong bases and the ionization of a weak base.

16.6 Weak Acids

Weak acids are introduced, and a method is presented for determining the pH of a weak acid solution. A simulation shows the correlation between the acid-ionization constant and the pH of a solution. You will also learn how the percent ionization of a weak acid depends on concentration.

16.7 Weak Bases

Weak bases are introduced, and a method presented for determining the pH of a weak base solution. You will use a simulation in this section to measure the pH of various solutions of weak bases. You will also learn how the percent ionization of a weak base depends on concentration.

16.8 Relationship Between Ka and Kb

A constant relationship between the ionization constants of conjugate pairs is presented.

16.9 Acid-Base Properties of Salt Solutions

Certain ions behave as acids or as bases in aqueous solution. A method for qualitatively determining the pH of salt solutions is introduced.

16.10 Acid-Base Behavior and Chemical Structure

Binary acids, oxyacids, and carboxylic acids are discussed, and their molecular structures related to trends in their acidity. The concepts of electronegativity and oxidation number are used to explain some of the trends.

16.11 Lewis Acids and Bases

Another perspective on acid-base theory is presented. The Lewis acid-base theory defines acids and bases broadly enough such that many more reactions can be categorized as acid-base chemistry than could be categorized under the Brønsted-Lowry theory. An animation presents the principles of the Lewis theory.


17 - Additional Aspects of Aqueous Equilibria

17.1. The Common-Ion Effect

Le Chtelier's principle is invoked to explain the common-ion effect. In this section you will view an animation of the common-ion effect at the molecular level.

17.2 Buffered Solutions

A buffer is a solution that prevents drastic pH change upon addition of acid or base. You will learn about the composition of a buffer and how to calculate the pH of a buffer using the Henderson-Hasselbalch equation. Two simulations allow you to determine the pH of various buffers before and after additions of strong acid or base.

17.3 Acid-Base Titrations

This section revisits acid-base titration and introduces pH titration curves. Strong acid–strong base and weak acid-strong base titrations are presented, and the differences between them are explored. An animation shows an acid-base titration, and you will learn to calculate the pH at any point in a titration.

17.4 Solubility Equilibria

In this section the concept of equilibrium is applied to the dissolution of sparingly soluble salts in water. The equilibrium constant for such a process is given a special subscript to denote the type of equilibrium to which it applies. The relationship between Ksp and solubility is explored, and you will learn to interconvert between them.

17.5 Factors that Affect Solubility

Solubility of a salt can be diminished by the addition of a common ion, or it can be enhanced by the removal of one of the ions via neutralization or complex-ion formation. The chemical factors that affect solubility are presented, and a movie demonstrates solubility being enhanced by the addition of acid.

17.6 Precipitation and Separation of Ions

Differences in solubility are exploited to separate metal ions. The metal ions are precipitated as insoluble salts by addition of an ion common to the salt. Salts of lowest solubility precipitate first, allowing the separation of different metal ions.

17.7 Qualitative Analysis for Metallic Elements

A qualitative analysis scheme is presented for determining the presence of common cations in solution.


18 - Chemistry of the Environment

18.1 Earth's Atmosphere

Earth's atmosphere is divided into four regions based on variation in temperature. Brief descriptions are given of each of these regions, with emphasis on the troposphere and the stratosphere.

18.2 The Outer Regions of the Atmosphere

The stratosphere is described in more detail, and two important photochemical processes are introduced: photodissociation and photoionization. An animation in this section details the formation and destruction of ozone in the upper atmosphere.

18.3 Ozone in the Upper Atmosphere

The discussion of ozone continues in this section. Both photochemical and thermochemical reactions are examined, and animations illustrate two different types of catalytic destruction of stratospheric ozone.

18.4 Chemistry of the Troposphere

The troposphere is the region of the atmosphere closest to Earth's surface. It consists almost entirely of nitrogen and oxygen. There are trace quantities of other gases in the troposphere as well, and their impact on weather and the environment will be presented. A video demonstration shows one of the more abundant trace atmospheric constituents, CO2, reacting with water to produce an acidic solution.

18.5 The World Ocean

The vast majority of Earth's water is seawater. This section explores the composition of seawater and how it is processed to provide a source of freshwater.

18.6 Freshwater

The tiny percentage of freshwater on the planet is supporting an increasing world population. Water quality and treatment are discussed in this section.


19 - Chemical Thermodynamics

19.1 Spontaneous Processes

Spontaneous processes are defined, and several examples given. Previously, we have seen that very exothermic reactions tend to be spontaneous, but there are many examples of spontaneous endothermic processes. The focus of this chapter is to bring together thermodynamic concepts for the purpose of reliably predicting the spontaneity of reactions and processes.

19.2 Entropy and the Second Law of Thermodynamics

The concept of disorder, first presented in connection with solution formation, is given a thermodynamic name, entropy. Entropy, symbolized S, is presented as a state function. You will see that processes leading to an increase in entropy tend to be spontaneous. Generalizations are made about the types of processes that lead to an increase in the entropy of a system. A movie illustrates the increasing disorder associated with solution formation.

19.3 The Molecular Interpretation of Entropy

Molecules can store energy in three forms of motion: translational, vibrational, and rotational. Each mode of motion is referred to as a degree of freedom. You will see that the more degrees of freedom a molecule has, the greater its entropy.

19.4 Calculation of Entropy Changes

Here you will learn to calculate the entropy change, , associated with a chemical reaction. You will see that thermodynamic tables list absolute entropies, in contrast to enthalpies that can only be tabulated as changes in enthalpy.

19.5 Gibbs Free Energy

Bringing together the concepts of enthalpy and entropy, both of which can serve as reasonably good predictors of spontaneity, we introduce the concept of free energy. For a process at constant pressure and constant temperature, a negative value for means that a reaction is spontaneous as written, under standard conditions.

19.6 Free Energy and Temperature

for a reaction changes with temperature. In this section you will learn how to determine at temperatures other than 25°C by assuming that and remain constant with changing temperature. You will see a demonstration of the formation of water from the combustion of hydrogen gas. An animation illustrates the decomposition reaction with which automobile air bags are inflated. A simulation allows you to explore the relationship between and temperature.

19.7 Free Energy and the Equilibrium Constant

In this section you will see that tells us essentially the same thing as K, namely whether an equilibrium lies to the right or to the left. And you will see that G gives us essentially the same information as Q, namely, in which direction the reaction has to proceed to achieve equilibrium.


20 - Electrochemistry

20.1 Oxidation-Reduction Reactions

Oxidation-reduction reactions are reviewed briefly. (They were introduced in eChapter 4.4.) You will view two animations of oxidation-reduction reactions in this section.

20.2 Balancing Oxidation-Reduction Reactions

The method of balancing reactions using half-reactions is presented. A video demonstration of an oxidation-reduction reaction is included in this section, and you will practice balancing more complex equations.

20.3 Voltaic Cells

Here the use of spontaneous oxidation-reduction reactions to produce electrical energy is introduced. The terms voltaic cell, cathode, anode, and salt bridge are introduced. You will see an animation of an electrochemical cell in this section.

20.4 Cell EMF

Electrons move from one species to another because of a difference in potential. In this section you will calculate the standard cell potential for a voltaic cell, using tabulated standard reduction potentials. The cell potential or electromotive force is a measure of the driving force for the reaction to occur. Two animations will show how the tabulated standard reduction potentials are determined—one at the macroscopic level and one at the molecular level.

20.5 Spontaneity of Redox Reactions

Cell potentials can be used to determine the spontaneity of an oxidation-reduction reaction. In this section you will see the connection between standard cell potential and standard change in free energy for a reaction. This section also contains a video demonstration of a spontaneous oxidation-reduction reaction.

20.6 Effect of Concentration of Cell EMF

In this section you will learn to calculate a cell potential under nonstandard conditions.

20.7 Batteries

Batteries are packaged voltaic cells that provide portable electrochemical energy. Several types of batteries are described and illustrated here.

20.8 Corrosion

Corrosion is described as an electrochemical process, and the reactions involved are given. Two methods of guarding against corrosion, specifically the rusting of iron, are presented.

20.9 Electrolysis

Essentially the opposite of what happens in a voltaic cell, electrolysis uses electrical energy to drive a nonspontaneous chemical reaction. Several examples of electrolysis are given, including an animation of the electrolysis of water and a video demonstration of electroplating.


21 - Nuclear Chemistry

21.1 Radioactivity

This section presents a review of alpha, beta, and gamma radiation and an animation of these three types of radiation being separated by electric and magnetic fields. Two other types of radioactivity are also introduced: positron emission and electron capture. You will learn to write nuclear equations in this section.

21.2 Patterns of Nuclear Stability

Empirical observations reveal certain patterns in the stability of nuclei. The neutron-to-proton ratio, which is illustrated by plotting the number of neutrons against the number of protons in stable nuclei, establishes a belt of stability. This is the area within which all stable nuclei are found. This section introduces certain magic numbers of neutrons and protons, which impart special stability to nuclei—just as there are certain numbers of electrons that impart chemical stability to atoms.

21.3 Nuclear Transmutations

Bombardment of nuclei with neutrons or with other nuclei leads to the production of heavy elements not found in nature. You will learn to write the equations and the shorthand for these processes.

21.4 Rates of Radioactive Decay

Radioactive decay is a first-order process. In this section you briefly review the equations that apply to first-order processes, an animation illustrates the phenomenon of half-life, and you learn how to use radioactivity-dating techniques. Finally, a simulation allows you to advance time and watch the mass of a radioactive substance fall.

21.5 Detection of Radioactivity

Photographic film and Geiger counters are used to detect radiation. Geiger counters are described and illustrated with a schematic diagram.

21.6 Energy Changes in Nuclear Reactions

In this section you will use the famous equation E = mc2 to determine the energy given off by a nuclear process where mass is not conserved.

21.7 Nuclear Fission

Fission, the process used in power plants and most nuclear weapons, is described in some detail. The subject of enrichment of uranium is discussed, and, relating to enrichment, a simulation allows you to compare the molecular speeds of various gases.

21.8 Nuclear Fusion

Fusion is described, and some pertinent equations are given.

21.9 Biological Effects of Radiation

The effects of ionizing radiation are described, and the different types of radiation are compared. Some of the units used to report doses of radiation are also given.


22 - Chemistry of the Nonmetals

22.1 General Concepts: Periodic Trends and Chemical Reactions

A brief review of the periodic trends discussed earlier is presented. Summaries are also given of the reaction types that have been encountered thus far.

22.2 Hydrogen

Hydrogen is a group in itself. The physical and chemical properties of hydrogen are summarized and a movie of the exothermic reaction between hydrogen and oxygen is shown.

22.3 Group 8A: The Noble Gases

Physical and chemical properties, along with sources and uses of the noble gases, are summarized.

22.4 Group 7A: The Halogens

Physical and chemical properties, sources, and uses of the halogens are summarized. This section contains two movies showing the behavior of the halogens.

22.5 Oxygen

The physical and chemical properties of oxygen are summarized, and information is given about the sources and abundance of the element. An animation illustrates the periodic trends in behavior of oxide compounds. There are also two movies illustrating reactions with oxygen and the acidic behavior of carbon dioxide in water.

22.6 The Other Group 6A Elements: S, Se, Te, and Po

Physical and chemical properties, along with abundance and sources of sulfur, are summarized. A movie illustrates the dehydrating capacity of the strong oxyacid, H2SO4.

22.7 Nitrogen

The physical and chemical properties of nitrogen are summarized. Abundance and sources of the element are also presented. This section contains a movie demonstrating the behavior of some oxides of nitrogen.

22.8 The Other Group 5A Elements: P, As, Sb, and Bi

Properties of phosphorus are summarized.

22.9 Carbon

Physical and chemical properties of carbon are summarized. The various allotropes of carbon are described, and their production and uses presented.

22.10 The Other Group 4A Elements: Si, Ge, Sn, and Pb

Properties, abundance, and uses of silicon are summarized.

22.11 Boron

A brief introduction to the properties of boron is given. This section contains an illustration of an unusual molecule, B2H6, in which hydrogen appears to form more than one bond.


23 - Metals and Metallurgy

23.1 Occurrence and Distribution of Metals

Most of the elements in the periodic table are metals, and several of these metals are useful and important to modern society. Most metals, however, do not occur in metallic form in nature. With their tendency to lose electrons, metals are typically found in solid inorganic compounds called minerals. Metallurgy is the processing of naturally occurring minerals to produce elemental metals. The basics of metallurgy are introduced in this section.

23.2 Pyrometallurgy

The oldest form of metallurgy, pyrometallurgy, is still used in the manufacture of steel. It involves the use of enormous furnaces and very high temperatures. You will see a demonstration of the thermite process, which will show how much of the heat necessary for the pyrometallurgy of iron comes from the reduction of iron itself.

23.3 Hydrometallurgy

Some metals can be concentrated and reduced using water or aqueous solutions. Several examples are presented, showing metals that can be concentrated and reduced using aqueous solutions.

23.4 Electrometallurgy

Production or purification of metals using electrolysis is called electrometallurgy. Sodium metal is produced by the electrolysis of molten sodium chloride. Copper metal is purified by electrolysis in which crude copper is used as the anode in an electrolytic cell.

23.5 Metallic Bonding

This section introduces the properties of metals and explains most of them with the electron-sea metallic bonding model. It also applies the molecular-orbital theory to metals in order to explain properties that cannot be explained using the electron-sea model.

23.6 Alloys

Many useful modern materials are metal alloys. This section presents the various types of alloys and gives examples of each. A table provides the composition of some familiar alloys.

23.7 Transition Metals

Trends in transition-metal properties are examined in much the same way that trends in main-group elements were examined in eChapter 7.3. Some of the trends are explained using electron configurations

23.8 Chemistry of Selected Transition Metals

Some interesting points of the chemistry of chromium, iron, and copper are presented. You will see a demonstration of an electrochemical reaction between iron and copper.


24 - Chemistry of Coordination Compounds

24.1 The Structures of Complexes

Compounds that contain complexes—a central metal ion bonded to surrounding molecules or ions—are called coordination compounds. The terms ligand, coordination sphere, and coordination number are introduced in this section. The geometries associated with various coordination numbers are shown. You will learn how to write formulas for metal complexes and how to determine the charge on the central metal ion.

24.2 Chelates

In this section you will see ligands that can attach to a metal ion with more than one bond. Such ligands are called polydentate ligands, or chelating agents. The rules of nomenclature for coordination compounds are introduced; you will practice deducing the name of a compound given its formula, and vice versa.

24.3 Isomerism

Compounds with the same formula may have different properties; these compounds have the same composition, but a different arrangement of atoms, and are called isomers. Animations in this section explore isomerism in general and a specific type of isomerism called optical activity.

24.4 Color and Magnetism

Some of the more interesting properties of coordination compounds are their colors and magnetic properties. Many coordination compounds are vividly colored, making them useful as pigments.

24.5 Crystal-Field Theory

Crystal-field theory presents a model for bonding in transition-metal complexes that accounts for our observation of color and magnetism in coordination compounds. You will see that the coordination of ligands to a metal ion causes the energies of the d orbitals on the metal ion change such that they are no longer degenerate. You will learn which ligands result in a large difference in d orbital energies and which result in a smaller difference. The magnitude of the difference influences both color and magnetic properties.


25- The Chemistry of Life: Organic and Biological Chemistry

25.1 A Look Back

Many of the concepts covered in earlier chapters are important for your understanding of organic and biological chemistry. A brief review is presented.

25.2 Introduction to Hydrocarbons

The class of compounds containing only carbon and hydrogen is presented. The terms saturated and unsaturated are introduced in the context of organic compounds. This section contains a simulation for determination of boiling points of hydrocarbons.

25.3 Alkanes

Nomenclature and combustion reactions of saturated hydrocarbons are discussed.

25.4 Unsaturated Hydrocarbons

Hydrocarbons containing one or more multiple bonds are introduced. Combustion and addition reactions are discussed.

25.5 Functional Groups: Alcohols and Ethers

This section introduces the term functional group for a group of atoms that impart properties to a molecule. Two functional groups, alcohols and ethers, are introduced. This section also makes use of the boiling point simulation to illustrate the effects of hydrogen bonding in alcohols.

25.6 Compounds with a Carbonyl Group

Several different functional groups contain a carbon-oxygen double bond, or carbonyl group. Several of these functional groups are introduced.

25.7 Chirality in Organic Chemistry

The concept of chirality is reviewed briefly in the context of organic molecules. An animation presents the concept of chirality; another illustrates the concept of optical activity.

25.8 Introduction to Biochemistry

This section serves as a very brief preview of the rest of the chapter. The functional groups in prior sections are integral to biologically important molecules.

25.9 Proteins

Proteins are introduced as one class of biopolymers. Amino acids are also introduced, and examples are given. The various structures and functions of proteins are discussed.

25.10 Carbohydrates

Compounds containing carbon, hydrogen, and oxygen are discussed. Simple sugars, disaccharides, and polysaccharides are discussed. In this section you will see how a tiny difference in the structure of a sugar can drastically change the properties of the polysaccharide derived from it.

25.11 Nucleic Acids

Nucleic acids, the carriers of genetic information, are also biopolymers. The building blocks of nucleotides are illustrated, and the function and replication of DNA are discussed briefly.