Correlations

Chemistry: The Central Science, Revised 8th Edition ©2002

Theodore L. Brown, H. Eugene LeMay, Jr., Bruce E. Bursten

Correlated with AP* Chemistry, May 1998, May 1999

ST = Student textbook pages

1. Structure of Matter
   1. Atomic theory and atomic structure
      1. Evidence for the atomic theory
         ST: 35–46
      2. Atomic masses; determination by chemical and physical means
         ST: 41–42, 74–86
      3. Atomic number and mass number; isotopes
         ST: 43, 228–229, 235, 808–806, 845
      4. Electron energy levels: atomic spectra, quantum numbers atomic orbitals
         ST: 185–209
      5. Periodic relationships including, for example, atomic radii, ionization energies, electron affinities, oxidation states
         ST: 10, 44–46, 51, 70–71, 121–122, 209–219, 226–259, 273, 282, 751
   2. Chemical bonding
      1. Binding forces
         1. Types: ionic, covalent, metallic hydrogen bonding, van der Waals (including London dispersion forces)
            ST: 261, 262, 263–272, 288–294 , 317–318, 333, 381–382, 395, 397–403, 476–477, 908–910
         2. Relationships to states, structure, and properties of matter
            ST: 261, 266–269, 277–282, 379–381, 421–426, 946–954
         3. Polarity of bonds, electronegativities
            ST: 273–278, 282, 315–317, 842
      2. Molecular models
         1. Lewis structures
            ST: 271–272, 278–282, 304
         2. Valence bond: hybridization of orbitals, resonance, sigma and pi bonds
            ST: 283–285, 317, 318–324, 325–331, 342, 842, 962, 977
         3. VSEPR
            ST: 304–314
      3. Geometry of molecules and ions, structural isomerism of simple organic molecules and coordination complexes; dipole moments of molecules; relation of properties to structure
         ST: 275–277, 303–314, 315, 318–324, 330–331, 360, 396–397, 854, 932, 940–944, 962, 966, 977
   3. Nuclear chemistry: nuclear equations, half-lives and radioactivity; chemical applications
      ST: 39, 406, 521–522, 805–839
   
   
2. States of Matter
   1. Gases
      1. Laws of ideal gases
         1. Equation of state for an ideal gas
            ST: 362–369, 373, 375
         2. Partial pressures
            ST: 369–373
      2. Kinetic-molecular theory
         1. Interpretation of ideal gas laws on the basis of this theory
            ST: 373–375
         2. Avogadro's hypothesis and the mole concept
            ST: 77–78, 361–362
         3. Dependence of kinetic energy of molecules on temperature
            ST: 373–375
         4. Deviations from ideal gas laws
            ST: 379–383
   2. Liquids and solids
      1. Liquids and solids from the kinetic-molecular viewpoint
         ST: 393–398
      2. Phase diagrams of one-component systems
         ST: 412–414
      3. Changes of state, including critical points and triple points
         ST: 11, 405–409, 412, 413
      4. Structure of solids; lattice energies
         ST: 5–6, 265–266, 267, 393–395
   3. Solutions
      1. Types of solutions and factors affecting solubility
         ST: 6, 105, 106–109, 130, 133, 469–475, 492, 488, 600, 601, 644–650
      2. Methods of expressing concentration (The use of normalities is not tested)
         ST: 482–486
      3. Raoult's law and colligative properties (nonvolatile solutes); osmosis
         ST: 486–495, 700, 701
      4. Non-ideal behavior (qualitative aspects)
         ST: 488–489
   
   
3. Reactions
   1. Reaction types
      1. Acid-base reactions; concepts of Arrhenius, Brønsted-Lowry, and Lewis; coordination complexes; amphoterism
         ST: 11–12, 114–120, 527, 529–532, 594–599, 627–631, 668–669, 929–959, 963
      2. Precipitation reactions
         ST: 109–113, 669–671
      3. Oxidation-reduction reactions
         1. Oxidation number
            ST: 121–122, 282, 751
         2. The role of the electron in oxidation-reduction
            ST: 120–121, 751
         3. Electrochemistry: electrolytic and galvanic cells; Faraday's laws; standard half-cell potentials; Nernst equation; prediction of the direction of redox reactions
            ST: 750–783, 785–786
   2. Stoichiometry
      1. Ionic and molecular species present in chemical systems: net ionic equations
         ST: 105–113, 131–136
      2. Balancing of equations, including those for redox reactions
         ST: 68–69, 86–90, 131, 155, 753–755
      3. Mass and volume relations with emphasis on the mole concept, including empirical formulas and limiting reactants
         ST: 14, 48, 77–83, 83–86, 91–94
   3. Equilibrium
      1. Concept of dynamic equilibrium, physical and chemical; Le Chatelier's principle; equilibrium constants
         ST: 108, 410, 475, 559, 562–568, 572–575, 576, 683, 641
      2. Quantitative treatment
         1. Equilibrium constants for gaseous reactions: Kp, Kc
            ST: 562–568
         2. Equilibrium constants for reactions in solution
            1. Constants for acids and bases; pK; pH
               ST: 601–605, 606–640, 646–648, 660–662
            2. Solubility product constants and their application to precipitation and the dissolution of slightly soluble compounds
               ST: 110–111, 659–663, 669–671
            3. Common ion effect; buffers; hydrolysis
               ST: 621, 629–630, 641–644, 644–650, 663–664
   4. Kinetics
      1. Concept of rate of reaction
         ST: 509, 510–532, 548, 713
      2. Use of differential rate laws to determine order of reaction and rate constant from experimental data
         ST: 509, 515–525
      3. Effect of temperature change on rates
         ST: 509, 525–532
      4. Energy of activation; the role of catalysts
         ST: 509–510, 539–548
      5. The relationship between the rate-determining step and a mechanism
         ST: 532–548
   5. Thermodynamics
      1. State functions
         ST: 152–153, 716
      2. First law: change in enthalpy; heat of formation; heat of reaction; Hess's law; heats of vaporization and fusion; calorimetry
         ST: 149–170, 289–291, 406, 713, 714
      3. Second law: entropy; free energy of formation; free energy of reaction; dependence of change in free energy on enthalpy and entropy changes
         ST: 719–737
      4. Relationship of change in free energy to equilibrium constants and electrode potentials
         ST: 737–741, 759, 771–773, 788–789
   
   
4. Descriptive Chemistry
   1. The authors state that they have always attempted to introduce students to descriptive chemistry by integrating examples throughout the text. Pertinent and relevant examples of "real" chemistry are woven into all of the chapters as a means to illustrate principles and applications. Chapters which more directly address the properties of elements and their compounds include Chapters 4 (Aqueous Reactions and Solution Stoichiometry), 7 (Periodic Properties of the Elements), 12 (Modern Materials), 18 (Chemistry of the Environnment), and 22–25 (Chemistry of the Nonmetals; Metals and Metallurgy; Chemistry of Coordination Compounds; the Chemistry of Life; Organic and Biological Chemistry). Descriptive chemistry is also incorporated in the end-of-chapter exercises. (Preface, p. xx) The Chemistry at Work and Chemistry and Life series of essays, through an emphasis on current world events, scientific discoveries, and medical breakthroughs, are intended to help students appreciate the real-world perspective of chemistry and the ways in which chemistry affects their lives. (Preface, p.xxi) See page xviii for a list of these essays.
      1. Chemical reactivity and products of chemical reactions
         ST: 68, 70–74, 145, 150, 157, 368–369, 510, 873, 982
      2. Relationships in the periodic table: horizontal, vertical, and diagonal with examples from alkali metals, alkaline earth metals, halogens, and the first series of transition elements
         ST: 10, 44–46, 55, 70–71, 110, 187, 213–214, 227–229, 243–248, 268, 605–606, 850–866, 912, 914–923
      3. Introduction to organic chemistry: hydrocarbons and functional groups (structure, nomenclature, chemical properties). Physical and chemical properties of simple organic compounds should also be included as exemplary material for the study of other areas such as bonding, equilibria involving weak acids, kinetics, colligative properties, and stoichiometric determinations of empirical and molecular formulas.
         ST: 48, 83–86, 260–301, 486–495, 508–557, 606–615, 875–876, 961–1011
   
   
5. Laboratory
   The first five chapters of the text introduce basic conjcepts, such as nomenclature. Stoichiometry, and thermochemistry that provide the necessary background for many of the laboratory experiments usually performed in general chemistry. (Preface, p.xix) The list of supplements to the text for the instructor (pp. xxii–xxiii) includes the following relevant to laboratory work:
   • Laboratory Experiments (Nelson/Kemp) (0-13-084101-3)
     This manual includes 41 finely tuned experiments chosen to introduce students to basic lab techniques and to illustrate core chemical principles. It contains pre-lab questions and detachable report sheets. This new edition has been revised to correlate more tightly with the text. Safety and disposal information has also been updated.
   • Annotated Instructor's Edition to Laboratory Experiments (0-13-084516-7)
     This AIE combines the full student lab manual with front and back appendices covering the proper disposal of chemical waste, safety instructions for the lab, descriptions of standard lab equipment and materials, answers to questions, and more.
   
   Chemical Calculations
   The following list summarizes types of problems either explicitly or implicitly included in the preceding material. Attention should be given to significant figures, precision of measured values, and the use of logarithmic and exponential relationships. Critical analysis of the reasonableness of results is to be encouraged.
   
   A series of essays in the text entitled Strategies in Chemistry provides advice to students on problem solving and "thinking like a chemist." (p.xix) See, for example, p. 79. See also the following pertaining to the problem-solving skills listed above:
   ST: 27, 54, 1001
   
   See the pages listed below for explanations in the text pertaining to the various problem types and for sample end-of-chapter exercises:
   1. Percentage composition
      ST: 76–77, 96–97
   2. Empirical and molecular formulas from experimental data
      ST: 48, 83–86, 99–100
   3. Molar masses from gas density, freezing-point, and boiling-point measurements
      ST: 79–81, 97–98, 367–368, 388, 399–400, 411, 412, 486–490, 502–503
   4. Gas laws, including the ideal gas law, Dalton's law, and Graham's law
      ST: 358–369, 373–375, 377–378, 385–387, 388–389
   5. Stoichiometric relations using the concept of the mole; titration calculations
      ST: 77–83, 97–98, 133–135, 140–141, 654
   6. Mole fractions; molar and molal solutions
      ST: 77–93, 97–100, 363, 367–368, 483–484, 485, 489–490, 502–503, 503–504
   7. Faraday's law of electrolysis
      ST: 244, 77 785–791, 800, 904
   8. Equilibrium constants and their applications, including their use for simultaneous equilibria
      ST: 562–568, 572–575, 585–588
   9. Standard electrode potentials and their use; Nernst equation
      ST: 762–778, 795–799
   10. Thermodynamic and thermochemical calculations
       ST: 144–176, 178–182, 713–741, 742–746
   11. Kinetics calculations
       ST: 508–547, 549–554