1. Chem. 1B – 10/27 Lecture

2. Announcements I Exam 2 –Thurs. (10/29) –Will cover: Ch. 16 (Titrations, Solubility, Complex Ions), Ch. 17 (all sections) –Similar Format to Exam 1 –Review of Exam Topics Today –Help Session This Afternoon (4:30 to 5:30 in Sequoia 452) Web Site –Link to Practice Exam 2 Plus Solutions –Have Quiz Solutions Through Quiz 6 Posted

3. Announcements II Lab –Experiment 6 Lab Report due Thursday –No Quiz next week Today’s Lecture –Thermodynamics Last example problem –Exam 2 Review –Electrochemistry (Ch. 18) Some review + basic concepts today

4. Chem 1B – Thermodynamics Chapter 17 – Equilibrium and Temperature A chemist has designed a catalyst allowing ethanol (C 2 H 5 OH) to be made from CO + H 2. The catalyst will only work at T > 150°C. At that temperature will the product still be favored? Determine the K at that temperature.

5. 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

6. 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 pK a 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 pK a of the conjugate weak acid affects the titration –Diprotic Acid/Base Titrations Know how to determine pK a1 + pK a2 from titration plots

7. Exam 2 Review Chapter 16 – Solubility Equilibria –Know how to set up K sp 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 increase by acid addition –Be able to predict if precipitation occurs

8. 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

9. 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

10. 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

11. 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  H, T and  S values –Know how  G relates to whether a process is spontaneous

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

13. Chapter 18 Electrochemistry Not on Exam 2 Electrochemical Reactions –Redox Reactions: A redox reaction is the coupling of an oxidation with a reduction These need to be coupled so that there is not net gain or loss of electrons –Definitions: Reduction: a reduction of the oxidation state (gain of electrons) Oxidation: an increase in the oxidation state (loss of electrons)

14. Chapter 18 Electrochemistry Electrochemical Reactions –Oxidation States: How do we determine these? Examples: H 2 O, NH 3, H 2 CO, MnO 4 -, SO 4 2- Note: examples with unusual oxidation states (Mn+7) are generally less stable (good as electrochemical reactants Electrochemical Reactions –Balancing Redox Reactions: 6 step method: 1)Assign oxidation states 2)Separate overall reaction into oxidation and reduction reactions

15. Chapter 18 Electrochemistry Electrochemical Reactions –Balancing Redox Reactions: 6 step method – cont. 3. Balance each half reaction with respect to mass in order a) mass all elements other than H, O, b) O by adding H 2 O, c) by adding H +, d) Add OH - to both side if in alkaline sol’n 4. Balance each half reaction for charge by adding electrons 5. Use common multiplier to get equal numbers of electrons for each half-reaction 6. Add each half reaction together to get net reaction without electrons as reactants or products Examples (unbalanced): AgNO 3 (aq) + Zn(s) ↔ Ag(s) + Zn(NO 3 ) 2 (aq) HClO(aq) + Fe 2+ (aq) ↔ Cl 2 (g) + Fe 3+ (aq) MnO 4 - (aq) + C 2 O 4 2- (aq) ↔ Mn 2+ (aq) + CO 2 (g)

16. Chapter 18 Electrochemistry Electrochemical Reactions – Different Forms –“Beaker” Reactions Products form along with heat (assuming  H < 0) Little control of reaction Products co-mingled (from reduction and oxidation) –Galvanic Cell Oxidation and reduction reactions may be divided into different parts (anode and cathode electrodes sometimes separated through two reaction cells) Reaction can be “harnessed” through voltage/power production Examples: batteries, pH measuring electrodes

17. Chapter 18 Electrochemistry Electrochemical Reactions – Different Forms –Electrolytic Cell In this type of cell, external electrical energy is used to force unfavorable reactions (e.g. 2H 2 O(l) ↔ 2H 2 (g) + O 2 (g)) to occur