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 Today’s Lecture 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 Chapter 17 – Entropy
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 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)

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)
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
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
GALVANIC CELL
voltmeter
Ag+ + e- → Ag(s)
Zn(s)
Ag(s) + –
AgNO3(aq)
ZnSO4(aq)
Zn(s) → Zn2+ + 2e-
Salt Bridge