I (amp) = C (coulumbs)/T (seconds)

Voltaic/galvanic cells - devices that use chemical reactions to produce an electric current

batteries are example of voltaic cells

Electrolysis- use of electric energy to effect a chemical change
Anode- the electrode at which oxidation occurs; negative ("an ox")
Cathode- electrode at which reduction occurs; positive ("red cat")
*the anode and cathode for electrolytic systems is the opposite those of a galvanic
Electrolytic Cell- electric current used to cause chemical change (reactant favored reactions)
Electromotive Force (emf)- drives the electron from the anode to the cathode
Standard Cell Potential-a quantitative measure of the tendency of reactants to proceed to products when all are in their standard states at 25 degrees Celsius

Using Standard Cell Potientals:
-all half reactions listed in the table can occur in either directions
-the more positive the reduction potential, the more likely the reduction of a substance will occur and the less likely the oxidation of the substance will occur
-when presented w/ 2 redox-active substances, the species reduced witll be the one w/ the more positive value for E red

Oxidation- loss of electron(s) by a species; oxidation number increases
Reduction- gain of electron(s) by a species; ox. number decreases

Direct redox reaction- both agents in direct contact. Single replacement. No current produced (no potential difference)
Indirect redox reaction- electrons transferred through external wire (from red. agent --> ox. agent)

Nernst Equation
E= Eo - (RT/nF)ln(Q)
E= potential of a cell under nonstandard condtions
Eo= standard cell potential
n=moles of electrons transferred
F= Faraday's constant (9.64853 x 10^4 J/V(mol))
R= Gas constant (8.3145 J/K(mol))
T= temperaute (K)
Q= Reaction quotient ([Product]/[Reactant])(M)

Eo and its relation to change of free energy

∆G= -nFEo
n= the number of moles(electrons) transfered
F= Faraday constant (9.65853 x 10^4 J/Vmol)
Eo= standard cell potential

Product favored reactions yield ∆Go < 0 and so Eo > 0
Reactant favored reactions yield ∆Go > 0 and so Eo < 0

Determining which way a reaction proceeds:
(#1) Cd --> Cd(2+) + 2e- or (#2) Cd(2+) +2e- --> Cd

and

(#1) Fe --> Fe(2+) + 2e- or (#2) Fe(2+) + 2e- --> Fe

Since Fe is the most bottom right on our chart of standard reduction potentials (in aqueous solution at 25 degrees Celsius), it is the most likely to be oxidized (and also the best reducing agent). Therefore, the reaction will start here and we use the first Fe half reaction

Fe --> Fe(2+) + 2e-

Then we pick the Cd half reaction that balances out the two electrons on the right side, and so we use the second Cd half reaction:

Fe + Cd(2+) + 2e- --> Fe(2+) + 2e- + Cd