The document discusses electrochemistry concepts including oxidation-reduction reactions, electrical chemical cells, galvanic cells, standard cell notation, standard cells, and the Nerst equation. Key points include: redox reactions involve separate oxidation and reduction half reactions; cells split these half reactions between an anode and cathode connected by a salt bridge; cell potential is measured in volts and relates to Gibbs free energy; and the Nerst equation allows predicting non-standard cell potentials based on reaction quotient concentrations.

1. By Logan Danielson

2. Electrochemistry
Oxidation Reduction chemistry (Electron Transfer)
Remember RedOx reactions must be balanced as two
separate reactions, the
Oxidation half reaction and the Reduction half reaction
Oxidation- involves the loss of an electron : Oil
Reduction- involves the gain of an electron: Rig

3. Electrical Chemical Cells
 Are split up to do the half reactions separately: an anode and a
cathode
 The anode handles the oxidation
 The cathode handles the reduction
 Electrons flow from the anode to the cathode
 The two half reactions are connected by the wire to allow the
electrons to flow and by a salt bridge
 The salt bridge prevents the build up of charge that would happen by
the flow of electrons to the cathode
 Cations from the salt bridge will flow into the cathode to dissipate the electron
charge build up from the presence of extra electrons in the cathode
 Anions from the salt bridge will flow into the anode to dissipate the electron
charge build up from the absence of extra electrons in the anode

4. Galvonic cell / Voltaic Cell
It does the work and produces energy
Cell potential – electro motive force
Electrical potential is measured in Volts, V
It is related in dG

5. Electrical Cell notation
/ Anode / Cathode
Metal connected to Anode | Anode reaction || Cathode reaction | Metal connect to Cathode
| = phase transition
|| = salt bridge
, = used to separate aqueous components
e.g.
Pt(s) | Fe2+(aq)(.1M),Fe3+(aq)(.2M) || Cu2+(aq) (.1M) | Cu(s)

6. Standard Cells
 All solutes at 1.0 M concentration
 All gases at 1 atm
 All solids present in pure form
 All standard potentials are based of off an electrical cell paired with Hydrogen
 Eocell (in volts)= Eocathode – Eoanode
 Eocell , Eocathode, and Eoanode are all written as reduction potentials
 Use to predict Eocell
 Used to predict spontaneously, if it was is a negative E then the reaction would happen in
the opposite direction
 Oxidizing agents: the agent is reduced, the strongest oxidizing agents have the highest
reduction potentials
 Reducing agents: the agent is oxidized, the strongest reducing agents have the lowest
reduction potentials

7. Electrical Work
Electrical work = charge * potential difference
Joules = coulombs * volts
Faraday constant, F, the magnitude of charge on one
mole of electrons
F= 9.65 *104 C/mol electron
w=-F * potential difference
wmax=-nFEcell
Because wmax=dG
dG=-nFEcell

8. Nerst equation
Eocell= .0592/n * log(K)
Ecell= Eocell - .0592/n * log(Q)
This allows us to predict non-standard reaction cell
potentials