1. Commercial Voltaic Cells. 3.7…or Applications of Voltaic Cells…

2. Commercial Voltaic Cells Problems with simple voltaic cells: – Voltage produced varies as the concentrations of the ions in the solution change – Current production is low Wants in a Voltaic Cell: – A large mass of reactants to produce current over a prolonged period – Cell that is rechargeable is nice. Returning the reagents to their original sites in the cell

3. Batteries Primary – cannot be returned to their original state by recharging so when reactants run out, battery is “dead” Secondary – “storage batteries” or “rechargeable batteries” – Reactions can be reversed

4. Primary Batteries LeClanche Dry Cell – Cells contain a moist paste of NH 4 Cl, ZnCl 2, and MnO 2. 2NH 4 + (aq) + 2e -  2NH 3 (g) + H 2 (g) Zn (s)  Zn +2 (aq) + 2e - – Gases can build up » Zn +2 (aq) + 2NH 3 (g) + 2Cl - (aq)  Zn(NH 3 ) 2 Cl 2 (s) » 2MnO 2 (s) +H 2 (g)  Mn 2 O 3 (s) +H 2 O(l) – Zn and ammonium can react (Acid and metal)… battery can leak

5. Primary Batteries

6. Alkaline Batteries – Generate current up to 50% longer than LeClanche dry cell – Use basic material instead of NH 4 Cl. NaOH or KOH

7. Secondary Batteries Recharging requires applying an electric current from an external source to restore the cell to its original state – Lead storage battery – Nickel-Cadmium (“Ni-cad”) batteries

8. Fuel Cells and Hybrid Cars Voltaic cells can only produce electric current until their reagents run out Fuel Cells – Reactants can be supplied continuously to the cell from an external resevoir Hydrogen-oxygen fuel cell – Hydrogen is pumped onto the anode of the cell – Oxygen is directed to the cathode O 2 (g) + 2H 2 O + 4e -  4OH - (aq) H 2 (g)  2H + (aq) + 2e -

9. Fuels Cells and Hybrid Cars

10. 2 halves separated by proton exchange membrane (PEM) – Protons (H + ) formed at the anode react with the hydroxide produced at the cathode… = H 2 O

11. Electromotive Force (EMF) Water only spontaneously flows one way in waterfall. Likewise, electrons only spontaneously flow one way in a redox reaction— from higher to lower potential energy

12. Electromotive Force (EMF) Electromotive Force – Difference in potential energy per electric charge Difference in potential energy of each side of the cell “force causing electrons to move” – cell potential and is designated E cell – Units are Volts (V) 1 V = 1 Joule/1 coulomb1 J = 1 V X 1 coulomb 1 coulomb is the quantity of charge that passes a point in an electric circuit when a current of 1 ampere flows for 1 second – 1 C = 1A x 1s

13. Measuring Standard Potentials To obtain standard potentials, cell voltages must be measured under standard conditions : – Reactant and products are present in their standard state – Solutes in aqueous solution have a concentration of 1.0M – Gaseous reactants or products have a pressure of 1.0 bar Cell potential under these conditions is called standard potential (E⁰ cell ) – And at 298K unless otherwise noted

14. Measuring Standard Potentials Set up standard half-cells and connect to Standard Hydrogen Electrode – Their values are referenced to a standard hydrogen electrode (SHE). – By definition, the reduction potential for hydrogen is 0 V: 2 H + (aq, 1M) + 2 e −  H 2 (g, 1 atm)

15. Measuring Standard Potentials When measuring, concentrate on 3 aspects of the cell: – The reaction that occurs – Direction of electron flow in the external circuit Electrons move from the electrode of higher potential energy to the one of lower potential energy – Cell potential

16. Standard Reduction Potentials The cell potential at standard conditions can be found through this equation: Because cell potential is based on the potential energy per unit of charge, it is an intensive property. E cell  = E red (cathode) − E red (anode) 

17. Standard Reduction Potentials – If we have values for E⁰ cathode and E⁰ anode we can calculate the standard potential (E⁰ cell ) – When the standard value of E⁰ cell is positive, the reaction as written is predicted to be product- favored at equilibrium If negative, predicted to be reactant-favored – If we know E⁰ cell and E⁰ anode or E⁰ cathode, we can calcuate the other… E cell  = E red (cathode) − E red (anode) 

18. Standard Reduction Potentials For the oxidation in this cell, E red = −0.76 V For the reduction, E red = +0.34 V

19. Standard Reduction Potentials = +0.34 V − (−0.76 V) = +1.10 V E cell  = E red  (cathode) − E red  (anode)

20. Tables of Standard Reduction Potentials Reduction potentials for many electrodes have been measured and tabulated

21. Tables of Standard Reduction Potentials 1.Reactions are written “oxidized form + electrons  reduced form” All potentials are for reduction reactions 2.The more positive the E⁰ cell value, the better the oxidizing ablility F 2 (g) is the best oxidizing agent in the table 3.The more negative the E⁰ cell value, the less likely the half-reaction will occur as a reduction, and the more likely the reverse half-reaction will occur (oxidation).

22. Tables of Standard Reduction Potentials 4.When a reaction is reversed (reduced form  oxidized form + electrons), the sign of E⁰ cell is reversed, but the magnitude is unaffected. 5.The reaction between any substance on the left (oxidizing agent) in the table with any substance lower than it on the right (reducing agent) is product-favored at equilibrium 6.The algebraic sign of the half-reaction reduction potential is the sign of the electrode when it is attached to the H 2 /H + standard cell

23. Tables of Standard Reduction Potentials 7.Electrochemical potentials depend on the nature of the reactants and products and their concentrations, not on the quantities of material used. Changing the stoichiometric coefficents for a half- reaction DO NOT change the value of E⁰ – Fe +3 + e -  Fe +2 E⁰ = +0.177 V – 2Fe +3 + 2e -  2Fe +2 E⁰ = +0.177 V

24. Standard Reduction Potentials Al (s) + 3Ag + (aq)  Al +3 (aq) + 3Ag (s) – Calculate the potential for that process – 2.459V E⁰ cell Al +3 (aq) + 3e -  Al (s) -1.66V Fe +2 (aq) + 2e -  Fe (s) -0.44V Ag + (aq) + e -  Ag (s) +0.799V

25. Oxidizing and Reducing Agents The strongest oxidizers have the most positive reduction potentials. – Very easily reduced The strongest reducers have the most negative reduction potentials. – Very easily oxidized

26. Oxidizing and Reducing Agents The greater the difference between the two, the greater the voltage of the cell.