1. Electrochemistry Hope you get a charge out of this one!!!!

2. 18.1 REDOX (It’s back) Remember this?!?!? Identify half reactions first, balance atoms first, then H +, H 2 O/OH -, and then electrons. Multiply reactions by integers to cancel out electrons. Practice Time!

3. 18.2 Galvanic Cells Redox reactions can be used to generate an electrical current. The electrons that are transferred during a redox reaction can be routed through a wire to perform work. Only when the circuit is complete and neutral will electrons flow.

4. Snazzy Animation

5. 18.2 Vocab Anode: Where oxidation occurs (the negative end) Cathode: Where reduction occurs (the positive end) Cell potential (E cell ): the driving force on the electrons (measured in volts). Galvanic cells and Voltaic cells are the exact same thing, just different names.

6. 18.3 Std. Reduction Potentials Based on the assignment of zero volts to the following reaction: The E cell = E oxd + E red Conditions are 1 atm, 1 M, and 298 K. Std. Reduction Potentials are written as reduction half-reactions. 2H + + 2e - H 2

7. Examples 2H + + 2e - H 2 E o = 0.00 V 2Hg 2+ + 2e - Hg 2 2+ E o = 0.91 V K + + e - K E o = -2.92 V

8. Important Tips Reduction potential is an intensive property: if the equation is multiplied by an integer (in the case of balancing) the potential stays the same. One of the reduction half-reactions must always be reversed – the E cell will be the difference between the two half reactions. A cell will always run spontaneously in the direction that produces a positive cell potential.

9. 18.4 Cell potential, electrical work, and free energy From the view of the system though… E is cell potential and –w is the work leaving the system. q is charge.

10. A note about charge (q) There are 96,485 Coulombs per mole of electrons, denoted F. q equals the moles of electrons times the charge per mole of electrons. q = nF where n is moles and F is Faraday constant.

11. Charge and Free Energy The maximum cell potential is directly related to the free energy difference between the reactants and the products in the cell.

12. 18.5 Dependence on Cell Potential on Concentration Think back to equilibrium…. Cell potential is dependent on concentration when the cell is considered an equilibrium reaction…the result is the Nernst Equation. Now, this guy stated the third law of thermodynamics, has a Nobel prize and helped establish physical chemistry…give the equation a break.

13. Nernst Equation

14. Electro Equilibrium When the cell reaction reaches equilibrium there is no longer any force driving the electrons (dead battery). Therefore, ΔG = 0 A cell will spontaneously discharge until Q = K and E cell = o.

15. 18.7 Corrosion The process of pure metals returning to their natural state, oxides (rusting). Most metals have reduction potentials less positive than oxygen, so when they are reversed and become oxidation reactions the E cell is positive. Most metals form an oxide coating that prevents further corrosion.

16. Examples GoldE cell = 1.50 V Oxygen E cell = 1.23 V Iron (II)E cell = -0.44 V

18. 18.8 Electrolysis This process is the reverse of a galvanic cell and requires a power supply. The E cell will always be NEGATIVE!!!! Cathode and anode switch and flow through the salt bridge is reversed. This process can be used to electroplate metals.

19. 18.8 Electrolysis We can measure the rate at which electrons are moving backwards, this current is called amperes. An amp is 1 coulomb of charge per second. How long must a current of 5.00 A be applied to a solution of Ag + to produce 10.5 g of silver metal.

20. Electrolysis of Water Have to add a bit of salt because pure water has so few ions. Produces H 2 and O 2.

21. Relative oxidizing abilities. The more positive the E o the greater the tendency to proceed in the direction indicated. When comparing, which ever species has the most positive E o will be the first reduced at a given voltage.