Home > Chemistry > Chemistry 1A Part Deux: Lecture 30 Notes

Chemistry 1A Part Deux: Lecture 30 Notes


This is also another lecture on electrochemistry.  Tough subject, learn what you can.


Electrochemical cells are used as portable power sources.  They don’t let the reaction go by itself, they split up the half cells, connect them though a wire with a switch.

Half Cell Reaction             Eo (volts)              dGo (kJ/mol)

Mg2+ + 2e- = Mg                -2.37                          +457

2H+ + 2e- = H2                       0                               0

F2 + 2e- = 2F-                        +2.87                      -544

These are equilibrium reactions.  Negative means the electron is unstable and wants to flow downhill.  Halogen species are quite stable.  Alkali metals are reactive.

Oxidation favored going up, reduction favored going down.  Halogens reduce, alkalis oxidize.

dEo cell = Eo cathode – Eo anode >0   Thus dGo cell <0


Two cells, one with a zinc electrode, one with a copper.  Connect with a wire and lightbulb.  Add in salt bridge to get ions moving.


dEo cell = 0.80 – -.76 = 1.56 V

Ni2+ +2e- = Ni    -.28V

Mg2+ + 2e- = Mg  -2.37 V

The species that is reduced is Ni2+.

Cu/Cu2+ // Ag+/Ag

d E cell o = Eo cathode – Eo anode

Oxidation is also called etching.  Reduction can be called plating.

Electrons flow down the potential difference and do work.

dG = -nFdE

F = 96500 C/mol

n is the number of electrons in mol/mol

dE is the cell potential in volts (1V = 1J/C)

dG = maximum work in kJ/mol

Calculating Cell Potential and Free Energy

Anode oxidation

Cu -2e- = Cu2+     +.34V

Cathode reduction

Ag+ + e- = Ag        +.80V

E cell = .46 V (height)

dGo = -2 96500 C/mol 0.46V = -89kJ/mol  (flow)

Overall Reaction: Cu + 2Ag+ = 2Ag + Cu2+

Concentrations cells have anodes and cathodes using the same reagent but different concentration.  Electrons will flow toward the more concentrated solution.  The more dilute solution is at a more energetic state.  Going from Cu of dilute to Cu2+ of concentrated.  Entropy drives this reaction.

Non Standard Conditions

dG = dGo + RTlnQ

Q = products/reactants

dG = -nFdE

dGo = -nFdEo

dE  = dEo – (.059log Q)/n

For the copper concentration cell,

dE = dEo – 0.59/n log [Cu2+ dilute/conc] = 0.03

Batteries are dry cells.  The alkaline dry cell is a primary battery.  Alkaline because there is hydroxide in it.  dE cell = 1.5V

The anode is zinc, which is oxidized.

Zn + 2OH- – 2e- = Zn(OH)2

The cathode is MnO2, which is reduced

1.5 WHr = 5.4 kJ

3 = 10.8

10 = 36

24 = 86

The Lead Acid Car Battery

Anode: Pb + SO4-2 -2e- = PbSO4

Cathode: PbO2 + SO4-2 + 4H+ + 2e- = PbSO4 + H2O

These reactions are reversed when a battery is being recharged.  6 cells in series.  Electrolyte of 37% H2SO4.  dE cell = 2.09.  Rechargeable, heavy.  Sulfuric acid is denser than water.

Nickel Metal Hydride Batteries

Anode: MH + OH- -e- = M + H2o

Cathode: NiO(OH) + H2O + e- = Ni(OH)2 + OH-

E cell = 1.3 volts

Litihium Ion Batteries

Anode: LiCoO2 -xe- = LiCoO2+

Cathode: xLi+ + 6C + xe- = LixC6

E cell = 3.6

MnO2 + H20 + e- = MnO(OH) + OH-


PE = U = mgh

PV = nRT

K = products/reactants

Electronegativity: NOF Cl Kr

dG = -nFdE

V = IR

c = wf

SALUTE  size activity location unit time equipment

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