Correlations

General Chemistry: An Integrated Approach, 4th Edition ©2005

John W. Hill and Ralph H. Petrucci

Correlated to: Advanced Placement* (AP*) Chemistry Course Description

SE = Student Edition
TE = Teacher Edition

I. Structure of Matter

A. Atomic theory and atomic structure
1. Evidence for the atomic theory	38–39; 64
2. Atomic masses; determination by chemical and physical means	42–43
3. Atomic number and mass number; isotopes	40–41; 45; 64; 265; 309; 800; 803
4. Electron energy levels: atomic spectra, quantum numbers atomic orbitals	285–288; 402–404; 404–410
5. Periodic relationships including, for example, atomic radii, ionization energies, electron affinities, oxidation states	315–317; 319; 317–319; 319–322; 322–323
B. Chemical bonding
1. Binding forces
2. Types: ionic, covalent, metallic hydrogen bonding, van der Waals (including London dispersion forces)	51–55; 340; 343–344; 347–353; 408–410; 464–465; 970
3. Relationships to states, structure, and properties of matter	348–349; 353–361; 819–821; 969–970
4. Polarity of bonds, electronegativities	349–353; 399–400
5. Molecular models
6. Lewis structures	348–349; 353–354; 357–360
7. Valence bond: hybridization of orbitals, resonance, sigma and pi bonds	402–404; 404; 409; 411–417
8. VSEPR	388–398; 409
9. Geometry of molecules and ions, structural isomerism of simple organic molecules and coordination complexes; dipole moments of molecules; relation of properties to structure	399; 400–402; 450
C. Nuclear chemistry: nuclear equations, half-lives and radioactivity; chemical applications	802–803; 804–808; 808–810; 821–823

II. States of Matter

A. Gases
1. Laws of ideal gases
—a. Equation of state for an ideal gas	185; 186; 188
—b. Partial pressures	193; 582–584; 586–587; 592–594
2. Kinetic-molecular theory
—a. Interpretation of ideal gas laws on the basis of this theory	171; 197–198; 199–203
—b. Avogadro's hypothesis and the mole concept	181; 181–184
—c. Dependence of kinetic energy of molecules on temperature	170–171; 197–203
—d. Deviations from ideal gas laws	188–190
B. Liquids and solids
1. Liquids and solids from the kinetic-molecular viewpoint	435–437; 437–440; 441
2. Phase diagrams of one-component systems	440–441
3. Changes of state, including critical points and triple points	441
4. Structure of solids; lattice energies	346; 464
C. Solutions
1. Types of solutions and factors affecting solubility	496–497; 498
2. Methods of expressing concentration (The use of normalities is not tested)	107; 484–491
3. Raoult's law and colligative properties (nonvolatile solutes); osmosis	502–506; 726
4. Non–ideal behavior (qualitative aspects)	493

III. Reactions

A. Reaction types
1. Acid-base reactions; concepts of Arrhenius, Brønsted-Lowry, and Lewis; coordination complexes; amphoterism	132; 653–656; 654
2. Precipitation reactions	138–140; 685; 689–690
3. Oxidation-reduction reactions	140; 143–144; 147; 148–151; 149; 749–755
—a. Oxidation number	140–142; 144–145
—b. The role of the electron in oxidation-reduction	150–151; 765–766
—c. Electrochemistry: electrolytic and galvanic cells; Faraday's laws; standard half-cell potentials; Nernst equation; prediction of the direction of redox reactions	126; 260; 421; 817
B. Stoichiometry
1. Ionic and molecular species present in chemical systems: net ionic equations	133; 136
2. Balancing of equations, including those for redox reactions	92–97
3. Mass and volume relations with emphasis on the mole concept, including empirical formulas and limiting reactants	47; 86–89; 101–103
C. Equilibrium
1. Concept of dynamic equilibrium, physical and chemical; Le Chatelier's principle; equilibrium constants	437; 575–577; 589–597; 645–646
2. Quantitative treatment
—a. Equilibrium constants for gaseous reactions: Kp, Kc	577–580; 582; 585–589; 732–740; 768; 769
—b. Equilibrium constants for reactions in solution
——i. Constants for acids and bases; pK; pH	589–594; 630–637
——ii. Solubility product constants and their application to precipitation and the dissolution of slightly soluble compounds	585–586
——iii. Common ion effect; buffers; hydrolysis	130; 591
D. Kinetics
1. Concept of rate of reaction	528; 531–532; 533
2. Use of experimental data and graphical analysis to determine reactant order, rate constants, and reaction rate laws	529; 532–535; 550–552
3. Effect of temperature change on rates	550–552
4. Energy of activation; the role of catalysts	559–563; 557–558
5. The relationship between the rate-determining step and a mechanism	557–560
E. Thermodynamics
1. State functions	221; 728–729
2. First law: change in enthalpy; heat of formation; heat of reaction; Hess's law; heats of vaporization and fusion; calorimetry	221; 230–247
3. Second law: entropy; free energy of formation; free energy of reaction; dependence of change in free energy on enthalpy and entropy changes	223–229; 713; 716–725
4. Relationship of change in free energy to equilibrium constants and electrode potentials	218–219; 728–729

IV. Descriptive Chemistry

1. Chemical reactivity and products of chemical reactions	4; 92–113; 135; 532; 539; 541; 543; 544; 735–737; 736
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	59–62; 373; 626; 651; 775; 838–845; 845–851; 942–943; 983
3. Introduction to organic chemistry: hydrocarbons and functional groups (structure, nomenclature, chemical properties).	57–59; 63–64; 132; 616; 622–623; 645–646; 653–656; 664

V. Laboratory

1. Making observations of chemical reactions and substances	N/A
2. Recording data	N/A
3. Calculating and interpreting results based on the quantitative data obtained	N/A
4. Communicating effectively the results of experimental work	N/A

Chemical Calculations

1. Percentage composition	103; 484
2. Empirical and molecular formulas from experimental data	47; 86–88; 88–89
3. Molar masses from gas density, freezing-point, and boiling-point measurements	81–83
4. Gas laws, including the ideal gas law, Dalton's law, and Graham's law	184; 190; 192; 193
5. Stoichiometric relations using the concept of the mole; titration calculations	78–83; 97; 98–112; 195; 229; 489
6. Mole fractions; molar and molal solutions	81–83; 488–489
7. Faraday's law of electrolysis	126; 260; 421; 781
8. Equilibrium constants and their applications, including their use for simultaneous equilibria	578; 582–583; 583–589
9. Standard electrode potentials and their use; Nernst equation	755; 757; 783–785
10. Thermodynamic and thermochemical calculations	230–237; 245–247
11. Kinetics calculations	170–171; 197–203; 215–216; 548