Chem. 1B – 10/25 Lecture

1. Chem. 1B – 10/25 Lecture

2. Announcements I • Exam 2: • Thursday (10/27) • Will Cover Titrations, Solubility, Complex Ions (from Ch. 16) + Chapter 17 (Thermodynamics) • Same Format as Exam 1 • Besides being able to do calculations (a big part of lectures), should know how to set up problems (e.g. calculation approach for titrations) • Help Session today 3:30 to 4:30 (replacing 2nd 30 min. of office hours) Sequoia 452 • Finishing Topic Review today

3. Announcements II • Lab/Quiz 7 • Quiz 7 next Monday and Tuesday (Experiment 8 + Electrochem basics) • Today’s Lecture • Review of Exam 2 Topics – Chapter 17 • Electrochemistry (Ch. 18 – Exam 3 material) • Review (Chapter 4.9 – Oxidation States, Redox Reactions) • Balancing Redox Reactions • Voltaic (or Galvanic) Cells

4. Exam 2 Review • Chapter 17 – Spontaneous Processes • Understand main concepts regarding spontaneous processes • 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)

5. Exam 2 Review • Chapter 17 – Entropy – cont. • Be able to predict the change in entropy for the surroundings based on the change in entropy for the system • 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

6. Exam 2 Review • 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 • 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)

7. Exam 2 Review • Chapter 17 – Gibbs Free Energy – cont. • Be able to calculate K from DGrxn° (or visa versa) • Know how temperature changes affect equilibrium shifts

8. Chapter 18 ElectrochemistryNot 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)

9. Chapter 18 Electrochemistry • Electrochemical Reactions • Oxidation States: • How do we determine these? • Examples: H2O, NH3, CaF2, H2CO, MnO4-, SO42- • Note: examples with unusual oxidation states (Mn+7) are generally less stable (good as electrochemical reactants) • Electrochemical Reactions • Balancing Redox Reactions: • 6 step method: • Assign oxidation states • Separate overall reaction into oxidation and reduction reactions

10. 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 H2O, 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

11. Chapter 18 Electrochemistry • Electrochemical Reactions • Balancing Redox Reactions – Cont. • Examples (unbalanced): AgNO3(aq) + Zn(s) ↔ Ag(s) + Zn(NO3)2(aq) HClO(aq) + Fe2+(aq) ↔ Cl2(g) + Fe3+(aq) MnO4- (aq) + C2O42-(aq) ↔ Mn2+(aq) + CO2(g)

12. Chapter 18 Electrochemistry • Electrochemical Reactions – Different Forms • “Beaker” Reactions • Products form along with heat (assuming DH < 0) • Little control of reaction • Products co-mingled (from reduction and oxidation) • Example: nail “rusts” (oxidation of Fe, reduction of O2) • Voltaic (Galvanic) Cells • Oxidation and reduction reactions may be divided into different parts (half-cells sometimes physically separated through two reaction cells) • Two electrodes are also needed • Reaction can be “harnessed” through voltage/power production • Examples: batteries, pH measuring electrodes

13. 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. 2H2O(l) ↔ 2H2(g) + O2(g)) to occur • Also requires two electrodes – but some differences from electrodes of voltaic cells • Examples: Production of Cl2 gas from NaCl(aq), production of H2 gas from water (above), instruments that measure degree of oxidation/reduction at specific voltages (analogous to spectrometers)

14. Chapter 18 Electrochemistry GALVANIC CELL • Voltaic Cells - Description of how example cell works • Reaction on anode = oxidation • Anode = Zn electrode (as the Eº for Zn2+ is less than for that for Ag+) • So, reaction on cathode must be reduction and involve Ag • Oxidation produces e-, so anode has (–) charge (galvanic cells only); current runs from cathode to anode • Salt bridge allows replenishment of ions as cations migrate to cathode and anions toward anodes voltmeter Ag+ + e- → Ag(s) Zn(s) Ag(s) + – AgNO3(aq) ZnSO4(aq) Zn(s) → Zn2+ + 2e- Salt Bridge