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)Dn. Worked example 15.6.
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, 638-639. 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+] = 10-pH. 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 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 Acid-Strong 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 DHfo 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 DSsys.
b. Briefly cover Standard Entropy, So. Emphasize Equation 18.6. Worked examples 18.2.
c. All. Qualitatively Predicting the Sign of DSosys (can also use DSrxn). 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.
4. All. Entropy Changes in the Universe. Emphasize importance of entropy of system AND entropy of surroundings.
a. Calculating DSsurr: derive and define based on DHsys (Equation 18.7).
b. 2nd Law of Thermodynamics: DSuniv (or DStot) = DSsys + DSsurr > 0. Worked example 18.4.
c. 3rd Law of Thermodynamics: Explain how S has a zero point thus we have exact S values, but DH and DG do not, thus we can never measure H or G values, only changes (DH and DG). Refer to Figure 18.5 for entropy of phase changes. Appendix 2 has more complete list of Thermodynamic data.
5. All. Predicting Spontaneity.
a. Define Gibbs Free Energy, DG. Show relationship between DSuniv and DG. DG = DH – TDS. DG < 0 for spontaneous process. Refer to Table 18.4 for sign combinations of DH and DS. Worked example 18.5.
b. Standard Free Energy, DGo. Define conditions for standard free energy. Compare criteria for DHfo and DGfo. Equation 18.12 (DGo = SnDGfo,prod – SnDGfo,react). Worked example 18.6.
c. Using DG and DGo to solve problems. DG determines whether or not a process is spontaneous. DGo is related to the magnitude of K and whether reaction is reactant- or product-favored. Like K, DGo 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.
6. All. Free Energy and Chemical Equilibrium. Review Q expressions. Equation 18.13 (G = DGo + RT lnQ). Equation 18.14 (DGo = -RT lnK). Worked example 18.8. Figure 18.6 illustrates the connection between DG (changes as a reaction proceeds; determines in which direction the rxn must go to reach equilibrium) and DGo (a constant value for a process). Table 18.5 relates K, DGo, and reactant-/product-favored. Worked examples 18.9 and 18.10.
7. Optional – Thermodynamics in Living Systems.
Review 3 ways to calculate DGorxn
1. DGof prod – DGof react
2. DGorxn = DHorxn - TDSorxn
3. DGorxn = -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.)