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Chem. 1B – 10/20 Lecture

This lecture provides a review of the topics covered in Exam 2, including titrations, solubility equilibria, and thermodynamics. It also includes example problems and practice questions. Join the help session on Tuesday for further clarification.

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Chem. 1B – 10/20 Lecture

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  1. Chem. 1B – 10/20 Lecture Updated in Announcement Slides No Quiz Week of 10/24

  2. Announcements I • Exam 2: • Week from today (10/27) • Will Cover Titrations, Solubility, Complex Ions (from Ch. 16) + Chapter 17 (Thermodynamics) • Help Session Tuesday in 3:30 to 4:30 Sequoia 452 • Mastering • Ch. 17 due Tuesday • Problem 3D answer is wrong

  3. Announcements II • Lab/Quiz 7 • Quiz 7 on Experiment 4 and 6 (pre-lab) + Ch. 17 • No Lab Quiz Next Week • Experiment 4 report due • Today’s Lecture • Thermodynamics • Effect of Temperature on Equilibrium Constants • More Practice Problems • Review of Exam 2 Topics

  4. Chem 1B – ThermodynamicsChapter 17 – Example Problems Catalysts can help energetically favorable reactions occur, but can not allow products to form if DGº of products is higher than reactants. Which of the following hydrocarbons can be produced by syngas (CO + H2) – assume H2O forms if needed? CH4 C2H6 CH3OH C2H2 How do we solve? Make balanced reactions and calculate DGº (or DHº and DSº) What are the best conditions for these reactions?

  5. Chem 1B – ThermodynamicsChapter 17 – Example Problems HI has a DGfº = 1.7 kJ/mol at 298 K Can it be formed from H2(g) + I2(s)? What is K for the reaction: H2(g) + I2(g) ↔ 2HI(g) if DGfº(I2(g)) = 19.3 kJ/mol Does increasing T favor reactants or products? Water is sprayed into a reaction flask at equilibrium and absorbs 99% of the HI but little of the other gases. Explain what this will do to DG.

  6. Chem 1B – ThermodynamicsChapter 17 – Equilibrium and Temperature We know DGº changes with temperature according to: DGº = DHº – TDSº (note: DHº and DSº may change with T – but generally not a lot) We also know that DGº = -RTlnK -RTlnK = DHº – TDSº or lnK = -DHº/RT + DSº/R A Plot of lnK vs. 1/T would give m (slope) = -DHº/R and b (y-intercept) = +DSº/R What would a positive slope in the above plot mean? What would a positive y-intercept mean?

  7. Chem 1B – ThermodynamicsChapter 17 – Equilibrium and Temperature A chemist has designed a catalyst allowing ethanol to be made from CO + H2. The catalyst will only work at T > 150°C. At that temperature will the product still be favored? Determine the K at that temperature.

  8. Chem 1B – ThermodynamicsChapter 17 – Example Questions Which of the following reactions leads to a decrease in entropy for the system? a) I2(s) ↔ I2(g) b) I2(s) ↔ I2(aq) c) I2(s) + H2(g) ↔ 2HI(g) d) 2I(g) ↔ I2(g) 2. Under what temperature regimes will the reaction: N2(g) + 3H2(g) ⇌ 2NH3(g) (DH° = -91.8 kJ/mol) be spontaneous? never b) high temperature c) low temperature d) always.

  9. Chem 1B – ThermodynamicsChapter 17 – Example Questions A reaction occurs which has DS < 0. We know that DSsurroundingsis ____________ (give sign) 4. Which of the following is likely to have the highest standard entropy? CH3OH(g) b) CH3OH(l) c) CH3OH(s) d) all are equal 5. The reaction H2(g) + I2(s) ↔ 2HI(g) has a DG° = +3.4 kJ mol-1. If a system starts with PH2 = 0.85 atm and PHI = 0.010 atm and T = 298K, DG =

  10. Chem 1B – ThermodynamicsChapter 17 – Example Questions Hydrazine (N2H4) is used to make rocket fuel and other products but it has positive DGf over all temperatures. A strategy to make it would be to: a) make from N2 and H2, but using higher temperatures b) make from N2 and H2, but use catalysts c) use a more stable reactant than H2 (such as CH4) d) use a less stable reactant than N2 (such as N2O) e) have the reaction also produce another unstable product (such as N3)

  11. Exam 2 Review • Chapter 16 - Titrations • In General: • Be able to calculate equivalence point volume or unknown concentration from other given information (e.g. 25 mL of [HX] requires 38.1 mL of 0.0830 M NaOH – find [HX]) • Recognize sharp vs. non-sharp titrations • Be able to determine titration type (e.g. diprotic weak acid titrated with strong base) from shape of titration curve • Be able to determine an appropriate indicator to use • Know what a titration error is and causes of titration errors • Strong Acid – Strong Base Titrations • Be able to calculate pH at any point in titration

  12. Exam 2 Review • Chapter 16 – Titrations – cont. • Weak Acid – Strong Base Titrations • Be able to calculate pH at any point in titration (particularly at ½ of equivalent volume and at equivalent volume) • Know how the pKa of the weak acid affects the titration • Weak Base – Strong Acid Titrations • Be able to calculate pH at any point in titration • Know how the pKa of the conjugate weak acid affects the titration • Diprotic Acid/Base Titrations • Know how to determine pKa1 + pKa2 from titration plots • Know what species are present at any point in a plot

  13. Exam 2 Review • Chapter 16 – Solubility Equilibria • Know how to set up Ksp reactions and equations for solubility reactions • Be able to calculate molar solubility in water • Be able to calculate molar solubility in a common ion (under “at equilibrium” assumptions) • Know qualitative effects of common ion addition • Know which salts can have solubility increased by acid addition • Be able to predict if precipitation occurs

  14. Exam 2 Review • Chapter 16 – Solubility Equilibria – cont. • Be able to calculate if an anion (or cation) can be added to selectively precipitate one of two cations (or anions) • Understand the basic methods used for qualitative analysis of ions • Chapter 16 – Complex Ion Formation • Understand basic nature of complex ion formation • Be able to solve equilibrium problems for complex ions under “at equilibrium” conditions • Know how complex ions affect solubility

  15. Exam 2 Review • Chapter 16 – General • Be able to combine equilibrium equations to predict importance of combined effects (example: combined solubility equilibrium + acid – base reaction to determine effect of acid on solubility or combined solubility equilibrium with complex ion formation) • Chapter 17 – Spontaneous Processes • Understand main concepts regarding spontaneous processes

  16. Exam 2 Review • Chapter 17 – Entropy • Understand basic concept of entropy • Be able to predict sign of entropy change for various processes (change in state, change in temperature, change in number of moles) • Know what state has an entropy of zero • Know the second law of thermodynamics (change in entropy for the universe) • Be able to predict the change in entropy for the surroundings based on the change in entropy for the system

  17. Exam 2 Review • Chapter 17 – Entropy – cont. • Be able to calculate the change in entropy for the surroundings based on the enthalpy change of the system and the temperature • Be able to calculate the standard change in entropy for a reaction using standard entropies of reactants and products • Chapter 17 – Gibbs Free Energy • Be able to calculate the Gibbs free energy change from DH, T and DS values • Know how DG relates to whether a process is spontaneous

  18. Exam 2 Review • Chapter 17 – Gibbs Free Energy – cont. • Be able to predict the temperature regime where a process is spontaneous from DH and DS information • Be able to calculate DG° for standard conditions from either DH°, T and DS° or from DGf° values • Know how DGrxn depends on reaction conditions (I will give equation: DGrxn = DGrxn° + RTlnQ) • Be able to calculate K from DGrxn° (or visa versa) • Know how temperature changes affect equilibrium shifts

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