1. Electrochemistry

2. What is Electrochem?
The relationship between chemical reactions and electricity.
Reactions create electricity
Electricity creates reactions

3. Electrochemical Cells
Two types
Cell produces electric current through a spontaneous reaction (galvanic aka voltaic)
Cell consumes electric current to push a non- spontaneous reaction (electrolytic cells)

4. Reactions creating electricity
Electricity = flow of electrons
Redox reactions = flow of electrons
A  e-  B = A  wire  B

5. Components:
A. Electrodes
Anode
Cathode
B. Solutions
C. Voltmeter
D. Salt Bridge

6. Oxidation always occurs at the anode.
Reduction always occurs at the cathode.
Zn half reaction lies further to the right
Why? Greater tendency to ionize than Cu.
Makes Zn negatively charged (anode) relative to Cu electrode (cathode).
Reaction driven by standard cell potentials.

7. Salt Bridges U-shaped tube with permeable stoppers.
Contains strong electrolytes (KCl, KNO3, etc.) usually suspended in gel.
Stops charge build-up in solutions for cells.

8. Standard Conditions Pressure = 1 atm for gaseous reactants.
Temperature = assume 25oC
Concentration= 1 M for each substance in solution

10. Electrochemical Cell Notation
Write oxidation half reaction on left and reduction half reaction on right.
Double vertical line (salt bridge) separates both half reactions
Different phases are separated by a vertical line. (separate same phases with a comma)

11. Electromotive Force Emf = cell voltage = Eocell
Driving force that pushes electrons away from the anode and towards cathode.
Eocell = Volt (V) = J/C
Coulomb = quantity of charge that passes a point in 1 sec when a current of 1 ampere flows.

12. Standard Cell Potentials
What is it? A measure of the overall tendency of the redox to occur. Eocell = Eocathode + Eoanode Eocell = positive for spontaneous reactions Eocell= negative for non-spontaneous reactions *All cell standards are reduction potentials. For oxidations, flip reaction and change the sign of potential.

13. Inert Electrodes
Pt and graphite are typically electrodes for gas phase and liquid reactions.
Ex. Standard Hydrogen Electrode (SHE) = uses Pt electrode.
SHE consists of Pt electrode in contact with M acid solution and H2 gas at 1 atm pressure.
H2(g)  2H+(aq) + 2e- Eored = 0.00 V

15. Standard Electrode Potential
The individual potential for each half reaction.
SHE used as standard for calculating potentials of other half reactions.
More negative charge = anode = more likely to undergo oxidation (think repelling e-s)
More positive charge = cathode = more likely to undergo reduction (think attracting e-s)

16. Strength of oxidizing agents
F2(g) + 2e-  2F-(aq)= most positive reduction potential
Easiest to reduce…why?
F2 = strongest oxidizing agent.

17. Strength of reducing agents
Li+ +e-  Li(s)
most negative Eored of V
Easiest to oxidize…why?
Strongest reducing agent
Wants to lose electrons, so reverse reaction occurs.
Li(s)  Li+ +e Eoox of V

18. Practice
Based on their reduction potentials, determine the best oxidizing agent and best reducing agent.
Au3+ + 3e-  Au(s) V
Br2(l) + 2e-  2Br-(aq) 1.07 V
Pb2+ + 2e-  Pb(s) V
Ni2+ + 2e-  Ni(s) V
Highest positive charge = Au = wants to gain electrons = best oxidizing agent.