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CHAPTER 16 (pages 776-792)
1. Oxidation and Reduction2. Galvanic Cells, Half Reactions (E°anode &
E°cathode)3. Standard Reduction Potential (E°)4. Nernst Equation, and the dependence of
Potential on Concentration5. Relationship between Equilibrium
Constant and Standard Potential
6. Driving Force, ΔG and ε1
REDOX REACTIONS
MnO2 + 4 HBr ⇋ MnBr2 + Br2 + 2 H2O
3 H2S + 2 NO3– + 2 H+ ⇋ 3 S + 2 NO + 4
H2O
2
OBSERVED REDOX PROCESSES
3
GALVANIC CELLS
4
INERT ELECTRODES
6
STANDARD REDUCTION POTENTIALS
7
8
MEASURING STANDARD POTENTIALS
9
CALCULATING STANDARD CELL POTENTIAL
Al(s) + NO3−
(aq) + 4 H+(aq)
⇋ Al3+
(aq) + NO(g) + 2 H2O(l)
10
ADDITIONAL EXAMPLE
Fe(s) + Mg2+(aq)
⇋ Fe2+
(aq) + Mg(s)
11
ox: Fe(s) Fe2+(aq) + 2 e− E = +0.45 V
red: Pb2+(aq) + 2 e− Pb(s) E = −0.13 V
tot: Pb2+(aq) + Fe(s) Fe2+(aq) + Pb(s) E = +0.32 V
ELECTROMOTIVE POTENTIAL
13
E°CELL, ΔG° AND K
Under standard state conditions, a reaction will spontaneously proceeds in the forward direction if:
– ΔG° < 1 (negative)– E° > 1 (positive)– K > 1
Design a voltaic cell with the following halfcells and complete the calculations:
Ag+ (aq) + 1e- Ag (s) Eo = 0.80 VPb2+ (aq) + 2e- Pb (s) Eo = -0.13 V
a. Calculate the Eocell
(potential at standard conditions)
b. Calculate Go.c. Calculate d. Calculate the Ecell if [Ag+] = 2.0 M and
[Pb2+] = 1.0 x 10-4 M.
Williams, spring 2009stop here
Design a voltaic cell with the following halfcells and complete the calculations:
Ag+ (aq) + 1e- Ag (s) Eo = 0.80 VPb2+ (aq) + 2e- Pb (s) Eo = -0.13 V
Calculate the Eocell
(potential at standard conditions)
Design a voltaic cell with the following halfcells and complete the calculations:
Ag+ (aq) + 1e- Ag (s) Eo = 0.80 VPb2+ (aq) + 2e- Pb (s) Eo = -0.13 V
Calculate Go.
Design a voltaic cell with the following halfcells and complete the calculations:
Ag+ (aq) + 1e- Ag (s) Eo = 0.80 VPb2+ (aq) + 2e- Pb (s) Eo = -0.13 V
Calculate
Design a voltaic cell with the following halfcells and complete the calculations:
Ag+ (aq) + 1e- Ag (s) Eo = 0.80 VPb2+ (aq) + 2e- Pb (s) Eo = -0.13 V
Calculate the Ecell if [Ag+] = 2.0 M and Pb2+] = 1.0 x 10-4 M.
OBJECTIVE 11.4: PROVIDE A THOROUGH OVERVIEW OF APPLICATIONS OF ELECTROCHEMICAL CELLS INCLUDING FUEL CELLS, CORROSION, AND OTHER TOPICS
AS TIME PERMITS.
23
CORROSION
• corrosion is the spontaneous oxidation of a metal by chemicals in the environment
• since many materials we use are active metals, corrosion can be a very big problem
RUSTING
• rust is hydrated iron(III) oxide• moisture must be present• electrolytes promote rusting• acids promote rusting– lower pH = lower E°red
Dry Cell Batteries
Lead – Acid Storage Battery
Biological Electrochemistry
Lithium Ion Battery