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Chemistry: Matter and Its Changes, 4th Edition

by Brady, James E., St. John's Univ., New York; Senese, Fred, Frostburg State University, Maryland

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Periodic Table of the Elements, FM-1
Table of Atomic Masses and Atomic Numbers, FM-2
Copyright, ii
Preface, iii
Acknowledgments, xv
To The Student, xviii
Facets Of Chemistry, xx
Brief Contents, xxi
Contents, xxii
Chapter 1. Atoms and Elements: The Building Blocks of Chemistry, xxviii
1.1. Chemistry is important for anyone studying the sciences, 1
1.2. The scientific method helps us build models of nature, 2
1.3. Properties of materials can be classified in different ways, 6
1.4. Materials are described by their properties, 8
1.5. Atoms of an element have properties in common, 13
1.6. Atoms are composed of subatomic particles, 20
1.7. The periodic table is used to organize and correlate facts, 25
Chapter 2. Compounds and Chemical Reactions, 40
2.1. Elements combine to form compounds, 41
2.2. Chemical equations describe what happens in chemical reactions, 45
2.3. Energy is an important part of chemical change, 47
2.4. Molecular compounds contain neutral particles called molecules, 49
2.5. Naming molecular compounds follows a system, 53
2.6. Ionic compounds are composed of charged particles called ions, 56
2.7. The formulas of many ionic compounds can be predicted, 58
2.8. Naming ionic compounds also follows a system, 63
2.9. Molecular and ionic compounds have characteristic properties, 69
Chapter 3. Measurement, 78
3.1. Measurements are quantitative observations, 79
3.2. Measurements always include units, 81
3.3. Measurements always contain some uncertainty, 90
3.4. Measurements are written using the significant figures convention, 93
3.5. Units can be converted using the factor-label method, 96
3.6. Density is a useful intensive property, 100
Chapter 4. The Mole: Connecting the Macroscopic and Molecular Worlds, 111
4.1. Use large-scale measurements to count tiny objects, 113
4.2. The mole conveniently links mass to number of atoms or molecules, 116
4.3. Chemical formulas relate amounts of substances in a compound, 121
4.4. Chemical formulas can be determined from experimental mass measurements, 125
4.5. Chemical equations link amounts of substances in a reaction, 133
4.6. Chemical equations cannot create or destroy atoms, 137
4.7. The reactant in shortest supply limits the amount of product, 139
4.8. The predicted amount of product is not always obtained experimentally, 142
TEST OF FACTS AND CONCEPTS Chapters 1–4, T-1
Chapter 5. Reactions Between Ions in Aqueous Solutions, 153
5.1. Special terminology applies to solutions, 154
5.2. Ionic compounds conduct electricity when dissolved in water, 157
5.3. Equations for ionic reactions can be written in different ways, 160
5.4. Reactions that produce precipitates can be predicted, 164
5.5. Acids and bases are classes of compounds with special properties, 167
5.6. Naming acids and bases follows a system, 172
5.7. Acids and bases are classified as strong or weak, 175
5.8. Neutralization occurs when acids and bases react, 179
5.9. Gases are formed in some metathesis reactions, 183
5.10. Predicting metathesis reactions—A summary, 185
5.11. The composition of a solution is described by its concentration, 189
5.12. Molarity is used for problems in solution stoichiometry, 195
5.13. Chemical analysis and titration are applications of solution stoichiometry, 201
Chapter 6. Oxidation–Reduction Reactions, 214
6.1. Oxidation–reduction reactions involve electron transfer, 215
6.2. The ion–electron method creates balanced net ionic equations for redox reactions, 224
6.3. Metals are oxidized when they react with acids, 229
6.4. A more active metal will displace a less active one from its compounds, 233
6.5. Molecular oxygen is a powerful oxidizing agent, 237
6.6. Redox reactions follow the same stoichiometric principles as other reactions, 241
Chapter 7. Energy and Chemical Change: Breaking and Making Bonds, 254
7.1. Energy is the ability to do work and supply heat, 255
7.2. Internal energy is the total energy of an object’s molecules, 260
7.3. Heat can be determined by measuring temperature changes, 263
7.4. Energy is absorbed or released when chemical bonds are broken or formed, 269
7.5. Heats of reaction are measured at constant volume or at constant pressure, 273
7.6. Thermochemical equations are chemical equations that quantitatively include heat, 280
7.7. Thermochemical equations can be combined because enthalpy is a state function, 282
7.8. Tabulated standard heats of reaction can be used to predict any heat of reaction using Hess’s law, 287
TEST OF FACTS AND CONCEPTS Chapters 5–7, T-3
Chapter 8. The Quantum Mechanical Atom, 302
8.1. Electromagnetic radiation can be described as a wave or as a stream of photons, 304
8.2. Atomic line spectra are experimental evidence that electrons in atoms have quantized energies, 311
8.3. Electron waves in atoms are called orbitals, 316
8.4. Electron spin affects the distribution of electrons among orbitals in atoms, 324
8.5. The ground state electron configuration is the lowest energy distribution of electrons among orbitals, 326
8.6. Electron configurations explain the structure of the periodic table, 328
8.7. Nodes in atomic orbitals affect their energies and their shapes, 334
8.8. Atomic properties correlate with an atom’s electron configuration, 337
Chapter 9. Chemical Bonding: General Concepts, 352
9.1. Electron transfer leads to the formation of ionic compounds, 353
9.2. Lewis symbols help keep track of valence electrons, 361
9.3. Covalent bonds are formed by electron sharing, 363
9.4. Carbon compounds illustrate the variety of structures possible with covalent bonds, 367
9.5. Covalent bonds can have partial charges at opposite ends, 371
9.6. The reactivities of metals and nonmetals can be related to their electronegativities, 374
9.7. Drawing Lewis structures is a necessary skill, 376
9.8. Formal charges help select correct Lewis structures, 381
9.9. Resonance applies when a single Lewis structure fails, 387
9.10. Both electrons in a coordinate covalent bond come from the same atom, 390
Chapter 10. Chemical Bonding and Molecular Structure, 400
10.1. Molecular shapes are built from five basic arrangements, 401
10.2. Molecular shapes are predicted using the VSEPR model, 404
10.3. Polar molecules are asymmetric, 413
10.4. Valence bond theory explains bonding as an overlap of orbitals, 417
10.5. Hybrid orbitals are used to explain experimental molecular geometries, 420
10.6. Hybrid orbitals can be used to explain multiple bonds, 431
10.7. Molecular orbital theory explains bonding as constructive interference of atomic orbitals, 437
10.8. Molecular orbital theory uses delocalized orbitals to describe molecules with resonance structures, 442
TEST OF FACTS AND CONCEPTS Chapters 8–10, T-5
Chapter 11. Properties of Gases, 450
11.1. Familiar properties of gases can be explained at the molecular level, 451
11.2. Pressure is a measured property of gases, 453
11.3. The gas laws summarize experimental observations, 459
11.4. The ideal gas law relates P, V, T, and the number of moles of gas, n, 468
11.5. In a mixture each gas exerts its own partial pressure, 475
11.6. Gas volumes are used in solving stoichiometry problems, 481
11.7. Effusion and diffusion in gases lead to Graham’s law, 484
11.8. The kinetic-molecular theory explains the gas laws, 486
11.9. Real gases don’t obey the ideal gas law perfectly, 491
Chapter 12. Intermolecular Attractions and the Properties of Liquids and Solids, 503
12.1. Gases, liquids, and solids differ because intermolecular forces depend on the distances between molecules, 504
12.2. Intermolecular attractions involve electrical charges, 505
12.3. Intermolecular forces and tightness of packing affect the properties of liquids and solids, 513
12.4. Changes of state lead to dynamic equilibria, 520
12.5. Vapor pressures of liquids and solids are controlled by temperature and intermolecular attractions, 523
12.6. Boiling occurs when a liquid’s vapor pressure equals atmospheric pressure, 526
12.7. Energy changes occur during changes of state, 528
12.8. Changes in a dynamic equilibrium can be analyzed using Le Châtelier’ s principle, 533
12.9. Phase diagrams graphically represent pressure–temperature relationships, 534
Chapter 13. Structures, Properties, and Applications of Solids, 547
13.1. Crystalline solids have an ordered internal structure, 548
13.2. X-ray diffraction is used to study crystal structures, 555
13.3. Physical properties are related to crystal types, 559
13.4. Band theory explains the electronic structures of solids, 561
13.5. Polymers are composed of many repeating molecular units, 564
13.6. Liquid crystals have properties of both liquids and crystals, 575
13.7. Modern ceramics have applications far beyond porcelain and pottery, 578
13.8. Nanotechnology deals with controlling structure at the molecular level, 586
Chapter 14. Solutions, 597
14.1. Substances mix spontaneously when there is no energy barrier to mixing, 599
14.2. Enthalpy of solution comes from unbalanced intermolecular attractions, 603
14.3. A substance’s solubility changes with temperature, 609
14.4. Gases become more soluble at higher pressures, 611
14.5. Molarity changes with temperature; molality, mass percentages, and mole fractions do not, 614
14.6. Substances have lower vapor pressures in solution, 618
14.7. Solutions have lower freezing points and higher boiling points than pure solvents, 622
14.8. Osmosis is flow of material through a semipermeable membrane due to unequal concentrations, 626
14.9. Ionic solutes affect colligative properties differently than nonionic solutes, 630
14.10. Colloids are molecular assemblies of submicrometer dimensions, suspended in a solvent, 633
TEST OF FACTS AND CONCEPTS Chapters 11–14, T-7
Chapter 15. Kinetics: The Study of Rates of Reaction, 645
15.1. The rate of a reaction is the change in reactant or product concentrations with time, 647
15.2. Five factors affect reaction rates, 647
15.3. Rates of reaction are measured by monitoring change in concentration over time, 650
15.4. Rate laws give reaction rate as a function of reactant concentrations, 654
15.5. Integrated rate laws give concentration as a function of time, 661
15.6. Reaction rate theories explain experimental rate laws in terms of molecular collisions, 669
15.7. Activation energies are measured by fitting experimental data to the Arrhenius equation, 674
15.8. Experimental rate laws can be used to support or reject proposed mechanisms for a reaction, 678
15.9. Catalysts change reaction rates by providing alternative paths between reactants and products, 683
Chapter 16. Chemical Equilibrium—General Concepts, 696
16.1. Dynamic equilibrium is achieved when the rates of two opposing processes are equal, 698
16.2. Closed systems reach the same equilibrium concentrations whether we start with reactants or with products, 699
16.3. A law relating equilibrium concentrations can be derived from the balanced chemical equation for a reaction, 700
16.4. Equilibrium laws for gaseous reactions can be written in terms of concentrations or pressures, 704
16.5. A large K means a product-rich equilibrium mixture; a small K means a reactant-rich mixture at equilibrium, 705
16.6. A simple expression relates K_P and K_c, 707
16.7. Heterogeneous equilibria involve reaction mixtures with more than one phase, 709
16.8. When a system at equilibrium is stressed, it reacts to relieve the stress, 711
16.9. Equilibrium concentrations can be used to predict equilibrium constants, and vice versa, 716
Chapter 17. Acids and Bases: A Second Look, 736
17.1. Brønsted–Lowry acids and bases exchange protons, 737
17.2. Strengths of Brønsted acids and bases follow periodic trends, 742
17.3. Lewis acids and bases involve coordinate covalent bonds, 751
17.4. Elements and their oxides demonstrate acid–base properties, 754
17.5. pH is a measure of the acidity of a solution, 758
17.6. Strong acids and bases are fully dissociated in solution, 766
Chapter 18. Equilibria in Solutions of Weak Acids and Bases, 774
18.1. Ionization constants can be defined for weak acids and bases, 775
18.2. Calculations can involve finding or using K_a and K_b, 781
18.3. Salt solutions are not neutral if the ions are weak acids or bases, 789
18.4. Simplifications fail for some equilibrium calculations, 796
18.5. Buffers enable the control of pH, 799
18.6. Polyprotic acids ionize in two or more steps, 808
18.7. Salts of polyprotic acids give basic solutions, 811
18.8. Acid–base titrations have sharp changes in pH at the equivalence point, 813
Chapter 19. Solubility and Simultaneous Equilibria, 831
19.1. An insoluble salt is in equilibrium with the solution around it, 832
19.2. Solubility equilibria of metal oxides and sulfides involve reactions with water, 844
19.3. Metal ions can be separated by selective precipitation, 847
19.4. Complex ions participate in equilibria in aqueous solutions, 852
TEST OF FACTS AND CONCEPTS Chapters 15–19, T-9
Chapter 20. Thermodynamics, 864
20.1. Internal energy can be transferred as heat or work, but it cannot be created or destroyed, 866
20.2. A spontaneous change is a change that continues without outside intervention, 870
20.3. Spontaneous processes tend to proceed from states of low probability to states of higher probability, 872
20.4. All spontaneous processes increase the total entropy of the universe, 878
20.5. The third law of thermodynamics makes experimental measurement of absolute entropies possible, 881
20.6. The standard free energy change, ΔG°, is ΔG at standard conditions, 884
20.7. ΔG is the maximum amount of work that can be done by a process, 887
20.8. ΔG is zero when a system is at equilibrium, 888
20.9. Equilibrium constants can be estimated from standard free energy changes, 895
20.10. Bond energies can be estimated from reaction enthalpy changes, 900
Chapter 21. Electrochemistry, 912
21.1. Galvanic cells use redox reactions to generate electricity, 913
21.2. Cell potentials can be related to reduction potentials, 919
21.3. Standard reduction potentials can predict spontaneous reactions, 924
21.4. Cell potentials are related to free energy changes, 930
21.5. Concentrations in a galvanic cell affect the cell potential, 933
21.6. Batteries are practical examples of galvanic cells, 937
21.7. Electrolysis uses electrical energy to cause chemical reactions, 945
21.8. Stoichiometry of electrochemical reactions involves electric current and time, 953
21.9. Electrolysis has many industrial applications, 956
Chapter 22. Nuclear Reactions and Their Role in Chemistry, 971
22.1. Mass and energy are conserved in all of their forms, 972
22.2. The energy required to break a nucleus into separate nucleons is called the nuclear binding energy, 973
22.3. Radioactivity is an emission of particles and/or electromagnetic radiation by unstable atomic nuclei, 975
22.4. Stable isotopes fall within the “band of stability” on a plot based on numbers of protons and neutrons, 981
22.5. Transmutation is the change of one isotope into another, 985
22.6. How is radiation measured?, 987
22.7. Radionuclides have many medical and analytical applications, 992
22.8. Nuclear fission is the breakup of a nucleus into two fragments of comparable size after capture of a slow neutron, 995
TEST OF FACTS AND CONCEPTS Chapters 20–22, T-11
Chapter 23. Metallurgy and the Properties of Metals and Metal Complexes, 1005
23.1. Metals are prepared from compounds by reduction, 1006
23.2. Metallurgy is the science and technology of metals, 1011
23.3. Metal compounds exhibit varying degrees of covalent bonding, 1016
23.4. Complex ions are formed by many metals, 1021
23.5. The nomenclature of metal complexes follows an extension of the rules developed earlier, 1026
23.6. Coordination number and structure are often related, 1028
23.7. Isomers of coordination complexes are compounds with the same formula but different structures, 1030
23.8. Bonding in transition metal complexes involves d orbitals, 1033
23.9. Metal ions serve critical functions in biological systems, 1041
Chapter 24. Some Chemical Properties of the Nonmetals and Metalloids, 1050
24.1. Metalloids and nonmetals are found as free elements and in compounds, 1051
24.2. The free elements have structures of varying complexity, 1054
24.3. Hydrogen forms compounds with most nonmetals and metalloids, 1059
24.4. Catenation occurs when atoms of the same element bond to each other, 1065
24.5. Oxygen combines with almost all nonmetals and metalloids, 1067
24.6. Nonmetals form a variety of oxoacids and oxoanions, 1073
24.7. Halogen compounds are formed by most nonmetals and metalloids, 1076
Chapter 25. Organic Compounds and Biochemicals, 1086
25.1. Organic chemistry is the study of carbon compounds, 1087
25.2. Hydrocarbons consist of only C and H atoms, 1092
25.3. Alcohols and ethers are organic derivatives of water, 1100
25.4. Amines are organic derivatives of ammonia, 1103
25.5. Organic compounds with carbonyl groups include aldehydes, ketones, and carboxylic acids, 1104
25.6. Most biochemicals are organic compounds, 1109
25.7. Carbohydrates include sugars, starch, and cellulose, 1109
25.8. Lipids comprise a family of water-insoluble compounds, 1112
25.9. Proteins are almost entirely polymers of amino acids, 1115
25.10. Nucleic acids carry our genetic information, 1119
Appendices, A-1
Appendix A. Electron Configurations of the Elements, A-1
Appendix B. Answers to Practice Exercises and Selected Review Problems, A-2
Appendix C. Tables of Selected Data, A-18
Glossary, G-1
Photo Credits, P-1
Index, I-1
Endpapers, EP-2

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