CHM 152 Topic Sequence Burdge/Overby, 1st Edition
14: Chemical Kinetics, 15: Chemical Equilibrium, 16: Acids and Bases, 17: Acid-Base and
Solubility Equilibria, 18: Entropy, Free Energy, and Equilibrium, 19: Electrochemistry,
20: Nuclear Chemistry
Worked Examples listed are merely suggested examples to show in class, work problems similar to,
and/or recommend that students be familiar with. Practice Problems A after each Worked example will
be similar problems; Practice Problems B after each Worked example will be an extension problem
requiring a more thorough understanding of the concept in order for students to solve.
Chapter 14 Chemical Kinetics (sections 1-8)
1. All. Reaction Rates. Overview, definition.
2. All. Collision Theory of Chemical Reactions.
a. Requirements for effective collisions (Figure 14.4 – two-page layout, p. 554); transition states;
energy profile diagrams (Figure 14.2); discuss determining endothermic or exothermic from graph;
define activation energy.
3. All. Measuring Reaction Progress and Expressing Reaction Rate (aka Rate expressions).
a. Average rates; Instantaneous rates; introduce k as a proportionality constant between rate and
concentration.
b. Qualitatively discuss the relationship between concentration and pressure using PV = nRT.
c. Stoichiometry and Reaction Rate: (Show Figure 12.1 from McMurry/Fay to show the relationship
between rates of chemical species in a reaction.) Worked examples 14.1 and 14.2.
4. All. Dependence of Reaction Rate on Reactant Concentration (aka Rate Laws). Worked example 14.3.
5. All. Dependence of Reactant Concentration on Time (aka Integrated Rate Laws). Worked examples 14.4 –
7 (emphasize that students will not have to generate a graph on a test, but they should be able to plot in
Excel for homework to determine order and unknown quantities based on graphs).
a. Qualitatively cover 0th order; show equation for linear plot.
6. All. Dependence of Reaction Rate on Temperature. Arrhenius Equation (Eqns: 14.10, 14.11). Worked
Examples 14.8, 14.9, 14.10 (again no generating graphs on tests, but students should be able to interpret
graphs and calculate Ea or k)
7. All. Reaction Mechanisms – elementary reactions; rate-determining step; determining rate laws when the
first step is slow (OR optional - determining rate laws when 2nd step) is slow. Worked example 14.11.
8. All. Catalysis. Figure 14.17. Define homogeneous vs. heterogeneous catalysts.
Key Skills two-page layout on First-Order Kinetics, pages 606-607.
End-of-chapter problems: Visualizing Chemistry (p. 595); 2, 4, 5, 9, 11, 12, 13, 14, 19, 21, 23, 25, 26 – 29, 31,
33, 37, 38, 39, 43, 47, 51, 53 – 56, 61, 63, 35 – 68, 77, 83, 85
Chapter 15 Chemical Equilibrium (sections 1-5)
1. All. The Concept of Equilibrium. Figure 15.1. Figure 15.2.
2. All. The Equilibrium Constant. Briefly explain the derivation of the equilibrium expression from rate laws.
Focus on equilibrium as equal rates of forward and reverse reactions, not equal concentrations.Show Table
15.1 to clarify that even with changing initial concentrations, Kc is a constant. Q vs K (Show Figure 15.4).
Worked Examples 15.1 and 15.2. Describe what the magnitude of equilibrium constants mean about an
equilibrium mixture.
3. All. Equilibrium Expressions. Heterogeneous expressions; Worked Example 15.3. Manipulating
equilibrium expressions (Table 15.2 summarizes – focus on 1st – reverse equation - and 4th – addition of
equations). Worked example 15.4. Writing Kp expressions. Worked example 15.5. Optional: showing
derivation of Kp = Kc(RT)n. Worked example 15.6.
Spring 2013
CHM 152
page 1 of 5
4. All. Using Equilibrium Expressions to Solve Problems. Using Q to determine direction of shift; Worked
example 15.7. Calculating equilibrium concentrations: setting up and solving ICE tables (perfect squares,
assuming small x, and quadratic formula); Worked Examples 15.8 – 15.10.
5. All. Factors that Affect Chemical Equilibrium. Le Chatelier’s Principle: Changes in concentration (Figure
15.5), Changes in volume/pressure, Changes in temperature, Catalysis (two-page layouts: 636-637, 638639. Worked examples 15.11 and 15.12.
Key Skills two-page layout on Equilibrium Problems, pages 654-655.
End-of-chapter problems: Visualizing Chemistry (pp. 645-646), 1 – 3, 6, 7, 8, 9, 11, 13 – 18, 21, 23, 25, 31, 35,
39, 41, 43, 45, 47, 49, 50 – 53, 55, 57, 59, 60, 61, 63, 65, 69, 73, 75, 84, 89
Chapter 16 Acid and Bases (sections 1-12)
1. All. Define Bronsted-Lowry acids and bases, emphasize recognition of conjugate pairs. Worked Examples
16.1 and 16.2. Explain to students that H+ is really H3O+.
2. All. Molecular Structure and Acid Strength. Hydrohalic acids: polarity of bonds determines bond
enthalpies which determines bond strength. Oxoacids: discuss trends of acids with different central atoms
and acids with the same central atom with different number of oxygen atoms. Optional: carboxylic acids
3. All. Acid-Base Properties of Water. Kw expression from autoionization of water. Kw =[H3O+][OH-] =
1.0x10-14 at 25oC.
4. All. The pH Scale. pH = -log [H3O+]. Define and emphasize significant figures rules for pH values
(decimal places determined by the sig. figs in concentration) – See page A-3 in Appendix 1. [H3O+] = 10pH
. pOH = -log [OH-]; [OH-] = 10-pOH; pH + pOH = 14.00. Worked Examples 16.5 – 16.8. Stress that
given one of these (H+, OH-, pH, pOH), the other three can be found.
5. All. Strong Acids and Bases. Define strong acids and bases. List and show equations for the 7 strong
acids (p. 668) and 8 strong bases (p. 669) with example pH calculations. Worked examples 16.9 – 16.12.
Note: Strong bases don’t react with water, they just dissociate producing hydroxide ions. Students need to
memorize the strong acids and strong bases. Stress that strong acids almost completely dissociate, not that
they have strong bonds. Emphasize exception of sulfuric acid: first proton is strong acid, second is weak
acid.
6. All. Weak Acids and Acid Ionization Constants (Ka). Emphasize small percent dissociation (thus small Ka
values) and equilibrium equation for weak acid solutions. Show generic (HA) formula for weak acids and
Ka expression. Table 16.7 shows example weak acids, structures and Ka values (more Ka and Kb values are
listed in Appendix 3). Two-page layout, pp. 674 – 675. Cover ICE tables to 1) calculate pH given initial
acid concentration and Ka, 2) calculate Ka given initial acid concentration and pH, and 3) calculate Ka
given initial acid concentration and percent ionization. Worked Examples 16. 13 – 16.15. Figure 16.5 to
show relationship between acid concentration and percent ionization.
a. Introduce approximation that x is small if Ka is small but must check approximation by 5% rule: If
[H+] / [HA]i x 100 < 5% approximation is good.
7. All. Weak Bases and Base Ionization Constants (Kb). Weak bases DO react with water (unlike strong
bases). Explain structure of amines (based on ammonia). Same concept as weak acids, except Kb is used.
Table 16.8 shows example formulas, structures, and Kb values. Worked examples 16.16 and 16.17.
8. All. Conjugate Acid-Base Pairs. Inverse relationship between strength of an acid or base and its
conjugate. Show derivation of KaKb = Kw Worked example 16.18.
9. Diprotic and Polyprotic acids. Show multiple dissociation steps for polyprotic acids and describe why K
values decrease with each proton removed. Qualitatively describe process to find pH for polyprotics. pH
is essentially determined from first ICE table. Emphasize exception of sulfuric acid: first proton is strong
acid, second is weak acid.
10. All. Acid-Base Properties of Salt Solutions. Focus on recognizing acidic and basic salts (contain
conjugate ions of weak acids and bases). Emphasize dissociation and hydrolysis equations. Also, label
ions as acidic, basic, or neutral. Strong acids plus weak bases make acidic salts (contain conjugate acid) so
they are acidic. Calculate pH for acidic and basic salts. Qualitatively discuss one example of a highly
Spring 2013
CHM 152
page 2 of 5
charged metal ion like Al3+ in figure 16.6. Worked examples 16.20 – 16.22. Compare K values of ions for
salts that have both acidic and basic ions.
11. Optional: Acid-Base Properties of Oxides and Hydroxides.
12. Briefly define Lewis acids and bases. Worked Example 16.23.
Key Skills two-page layout on Salt Hydrolysis, pp. 710 – 711.
End-of-chapter problems: 2, 4, 5, 9, 11, 14, 18, 20, 21, 24, 25, 26, 29, 32, 34, 35, 39, 41, 45, 51, 54, 57, 59, 61, 65,
72, 73, 77, 80, 83, 93, 95, 96, 97, 100, 101
Chapter 17 Acid-Base Equilibria and Solubility Equilibria (sections 1-6)
1. All. Common Ion Effect. Really drive home the concept. Worked example 17.1.
2. All. Buffer Solutions. Focus on the definition and concept of a buffer. Do calculations on pH of buffer
solutions and what happens if small amounts of strong acids or strong bases are added. Optional –
derivation of Henderson-Hasselbalch equation (Equation 17.1). Show example using equation. Worked
example 17.2. Two-page layout on buffers, pp. 718 – 719. Briefly cover how to make buffer solutions to
a certain pH (reference the buffer solutions used to calibrate pH probes in lab).
3. All. Acid-Base Titrations.
a. Strong Acid-Strong Base: reference Figure 17.3 and Table 17.1. Each step can be thought of as a
limiting reagent problem. What substance and how much is in excess determine the pH.
b. Weak Acid-Strong Base: reference Figure 17.4 and Table 17.2. Work through the 4
regions/points in a titration, focusing on the buffer zone and equivalence point.
c. Qualitatively discuss Strong Acid-Weak Base titrations and how they compare to Weak AcidStrong Base, reference Figure 17.5. Optional – quantitative calculation of pH values.
d. Acid-Base Indicators: Emphasize the difference between equivalence points and end points.
Qualitatively discuss indicators so students understand that phenolphthalein is not the only option
(reference Table 17.3).
4. All. Solubility Equilibria. Review saturated vs unsaturated solutions. Review Solubility rules; soluble
salts (in Chapter 16) vs. insoluble salts (Chapter 17). Focus on Ksp describing salt dissociating in water
equation.
a. Calculate Ksp given solubility (g/L) or molar solubility (mol/L) and vice versa. Refer to Table 17.4
for common Ksp values (more values in Appendix 4). Worked examples 17.7, 17.8.
b. Determining whether precipitation will occur using Qsp vs. Ksp values. Worked example 17.9.
5. All. Factors Affecting Solubility. Common Ion Effect (decreases solubility – relate to Le Chatelier’s
Principle). Two-page layout, pp. 738 – 739. Worked example 17.9. Qualitatively discuss solubility based
on pH of solutions (e.g., Mg(OH)2). Define Kf and complex ion formation (Figure 17.9) but no
calculations.
6. Optional - Separation of Ions Using Differences in Solubility. Figure 17.12 shows flowchart for separation
methods – helpful as reference for Qualitative Analysis experiment.
Key Skills two-page layout on Buffers: pp. 760 – 761.
End-of-chapter problems: 5, 6, 9, 11, 12, 17, 18, 21, 24, 25, 29, 31, 34, 35, (36 - weak base-strong acid titration),
42, 45, 51, 52, 54, 57, 65, 67, 68, 113
Chapter 18 Thermodynamics (sections 1-7) – Refer to page 372 and section 10.6 for review.
1. All. Define spontaneous vs nonspontaneous (refer to Table 18.1). Review endothermic vs. exothermic
reactions; predicting and calculating signs of enthalpy based on types of reactions/phase changes. Review
ways to calculate enthalpies (primarily using Hfo values). Review system vs. surroundings.
2. Entropy: Qualitatively define. Discuss number of possible arrangements to determine entropy. Optional
– mathematical overview of determining entropy.
3. Entropy Changes in a System. Two-page layout, pp. 772 – 773.
a. Optional – Calculating Ssys.
Spring 2013
CHM 152
page 3 of 5
4.
5.
6.
7.
b. Briefly cover Standard Entropy, So. Emphasize Equation 18.6. Worked examples 18.2.
c. All. Qualitatively Predicting the Sign of Sosys (can also use Srxn). Discuss phase changes,
temperature changes, and different numbers of gas molecules. Discuss the dissolution process –
not always increasing entropy (refer to Thermodynamics experiment). Worked example 18.3.
All. Entropy Changes in the Universe. Emphasize importance of entropy of system AND entropy of
surroundings.
a. Calculating Ssurr: derive and define based on Hsys (Equation 18.7).
b. 2nd Law of Thermodynamics: Suniv (or Stot) = Ssys + Ssurr > 0. Worked example 18.4.
c. 3rd Law of Thermodynamics: Explain how S has a zero point thus we have exact S values, but 
and G do not, thus we can never measure H or G values, only changes (H and G). Refer to
Figure 18.5 for entropy of phase changes. Appendix 2 has more complete list of Thermodynamic
data.
All. Predicting Spontaneity.
a. Define Gibbs Free Energy, G. Show relationship between Suniv and G. G = H – TS. G
< 0 for spontaneous process. Refer to Table 18.4 for sign combinations of H and S. Worked
example 18.5.
b. Standard Free Energy, Go. Define conditions for standard free energy. Compare criteria for Hfo
and Gfo. Equation 18.12 (Go = nGfo,prod – nGfo,react). Worked example 18.6.
c. Using G and Go to solve problems. G determines whether or not a process is spontaneous.
Go is related to the magnitude of K and whether reaction is reactant- or product-favored. Like K,
Go changes with changing temperatures. Solve for temperatures at which processes become
spontaneous (or non-spontaneous) as well as temperatures of phase changes). Worked example
18.7.
All. Free Energy and Chemical Equilibrium. Review Q expressions. Equation 18.13 (G = Go + RT
lnQ). Equation 18.14 (Go = -RT lnK). Worked example 18.8. Figure 18.6 illustrates the connection
between G (changes as a reaction proceeds; determines in which direction the rxn must go to reach
equilibrium) and Go (a constant value for a process). Table 18.5 relates K, Go, and reactant-/productfavored. Worked examples 18.9 and 18.10.
Optional – Thermodynamics in Living Systems.
Review 3 ways to calculate Gorxn
1. Gof prod – Gof react
2. Gorxn = Horxn - TSorxn
3. Gorxn = -RT lnK
Key Skills two-page layout on Free Energy and Equilibrium, pp. 802 – 203.
End-of-chapter problems: 2, 5, 9, 11, 15, 17, 21, 23, 27, 31, 33, 37, 41, 43, 55, 59, 61, 69, 75, 77, 95, 101, 103
Chapter 19 Electrochemistry (sections 1-8)
1. All. Balancing Redox Reactions. Review oxidation states, oxidation, reduction, oxidizing agents, and
reducing agents (from Chapter 9). Optional - Balance redox reactions (simple or in acid/base
solutions???). Worked example 19.1.
2. All. Galvanic Cells. Introduce galvanic (or voltaic) cells with simple rxns using metal electrodes in
corresponding solutions. Draw cell diagrams and discuss everything: anode ½ rxn, electrons travel via
wire (not through salt bridge), cathode ½ rxn, charge buildup in solutions thus salt bridge function and ion
flow, electrode mass changes, balancing electrons in ½ reactions. Show short-hand notation for galvanic
cells. (Two-page layout on Galvanic cells, pp. 810 – 811.)
3. All. Standard Reduction Potentials. Discuss SHE (Standard Hydrogen Electrode). Show Figures 19.3 and
19.4 as examples of galvanic cells with non-solid half-reaction. Review standard conditions (Eo). This
book does not use the term electromotive force, EMF, but this is where we get the letter E for cell potential
Spring 2013
CHM 152
page 4 of 5
4.
5.
6.
7.
8.
from. Calculate Eocell using ½ red potentials in Table 19.1. Use examples where ½ rxns are not just simple
M  M+2,some metals can be reduced to either ions or solids. Compare oxidizing and reducing agent
strength. Worked example 19.2.
All. Spontaneity of Redox Reactions Under Standard-State Conditions. (Optional: describing relationship
between work and cell potential.) G = -nFEcell, Go = -RT ln K, Eo = (RT/nF) ln K, Eo = (0.0592/n) log
K (students can use any variation they want to, but only the first equation is given on the equation sheet).
Worked examples 19.4, 19.5.
All. Spontaneity of Redox Reactions Under Conditions Other Than Standard State (E vs Eo). Nernst
equation: Show both variations: E = Eo - (RT/nF) ln Q (Equation 19.6); E = Eo - (0.0592/n) log Q (at
25oC) (Equation 19.7), but only the first one is on the equation sheet. Worked example 19.6. (Optional:
concentration cells: same materials with different concentrations of solutions.) Worked example 19.7.
All. Batteries – only show examples (show pictures and relate components to galvanic cells); Figures 19.6
– 19.9.
Electrolysis. Qualitative discussion of electrolysis. Quantitative applications of electrolysis (refer to
Figure 19.12). Worked example 19.8 (Key Skills two-page layout on Electrolysis of Metals, pp. 848 –
849).
Optional – Corrosion.
End-of-chapter problems: 3-6, 8, 9, 11,13, 15, 17, 18, 21, 23, 25, 29, 31, 33, 37, 42, 45, 47, 51, 53, 57, 72, 84, 125
Chapter 20 Nuclear (sections 1-8) Introduction: Cover notations and review of nuclear particles and notations
(atomic number, mass number, and isotopes).
1. All. Nuclei and Nuclear Reactions. Discuss radioactivitiy (spontaneous decay) and nuclear transmutation
(bombardment of nuclei). Introduce symbols for subatomic particles. Discuss balancing nuclear
equations. Worked Example 20.1.
2. Nuclear Stability. Optional – calculations of density of nucleus. Cover patterns of stability, Figure 20.1.
Brief coverage of band of stability. Above the band is often beta decay, below is positron decay or ecapture, higher mass often alpha decay. Everything above Bi is radioactive. Optional - define Nuclear
Binding Energy and NBE calculations.
3. All. Natural Radioactivity. Decay series for Uranium-238. Kinetics of Radioactive Decay (follow firstorder kinetics). Qualitatively discuss radioactive dating (using “atomic clocks”). Worked examples 20.3
and 20.4.
4. All. Nuclear Transmutation. Optional – parenthetical notation used to describe transmutation equations
(p. 862). Figure 20.5 shows cyclotron particle accelerator. Table 20.3 is interesting for students to see
formation of elements 93 – 118.
5. Optional – Nuclear Fission. Fission is process of a heavy nucleus being bombarded by a neutron to form
smaller nuclei and one or more neutrons. Two-page layout, pp. 866-867. Define nuclear chain reaction,
critical mass, and schematics of atomic bomb (Figure 20.10) and nuclear power plant (Figure 20.11).
6. Optional – Nuclear Fusion. Fusion is process of combining small nuclei into larger ones.
7. Optional – Uses of Isotopes: medicine.
8. Optional – Biological Effects of Radiation. Define units of curie, rad, and rem. Table 20.5 is interesting
for students to see.
As time may be short, sections 1 – 3 and defining reaction types, balancing equations, and identifying missing
particles in equations are most important.
End-of-chapter problems: 1-5, 7, 15, 25, 27, 29, 31, 65, 69, 71, 101
Spring 2013
CHM 152
page 5 of 5