1. Voltaic/Galvanic Cells

2. Voltaic Cells In spontaneous oxidation-reduction (redox) reactions, electrons are transferred and energy is released.

3. Voltaic Cells We can use that energy to do work if we make the electrons flow through an external device. We call such a setup a voltaic cell.

4. Voltaic Cells p. 836 A typical cell looks like this. The oxidation occurs at the anode. The reduction occurs at the cathode.

5. Voltaic Cells Once even one electron flows from the anode to the cathode, the charges in each beaker would not be balanced and the flow of electrons would stop. OH NO!!!!!!!!!!!!!!!!!!!!!

6. Therefore, we use a salt bridge, usually a U- shaped tube that contains a salt solution, to keep the charges balanced. –Cations move toward the cathode. –Anions move toward the anode.

7. Let’s label a little! Consider the following eqn: Zn +2 (aq) + Cu (s) ↔ Cu 2+ (aq) + Zn (s) Label the following: – Anode – Cathode – Direction of electron flow

8. Voltaic Cells In the cell, electrons leave the anode and flow through the wire to the cathode. As the electrons leave the anode, the cations formed (LEFT BEHIND) dissolve into the solution in the anode compartment.

9. Voltaic Cells As the electrons reach the cathode, EXISTING cations in the cathode are attracted to the now extra negative cathode. The electrons are taken by the cation, and the neutral metal is deposited on the cathode. Watch This

10. Let’s make a mini-voltaic cell. Get a Demonstration of a Voltaic Cell Sheet. Go Online to link.

11. Voltaic Cell Diagram Representation of the overall reaction in the electrochemical cell. The chemicals involved are what are actually reacting during the reduction and oxidation reactions. Makes it easier to see what is being oxidized and what is being reduced, focus is on reactions that create the cell potential.

12. Cell Diagrams The anode is always placed on the left side. The cathode is placed on the right side. The salt bridge is represented by double vertical lines (||).

13. Back to your Diagram? Write the cell diagram for the redox reaction taking place under the diagram. Use the appropriate abbreviations.

14. Cell Potential Electrons only spontaneously flow one way in a redox reaction—from higher to lower potential energy, creating a potential difference. We quantify this amount and call it Cell Potential (E cell ) E o cell = E o right (cathode) – E o left (anode)

15. Example 1. Split the reaction into half reactions and determine their E o value. USE PURPLE SHEET! Indicate which would be the anode and cathode. 2. Construct a cell diagram for the reactions. 3. Determine the E o cell for the cell formed by each. Consider the following two reactions: a) Cu 2+ (aq) + Ba (s) → Cu (s) + Ba 2+ (aq) b) Al (s) + Sn 2+ (aq) → Al 3+ (aq) + Sn (s) Reduction Value is opposite/flip/inverse of Oxidation Value.

16. Answers a) Ba 2+ (aq) + 2e- → Ba (s) E o = -2.92 V Anode Cu 2+(aq) + 2e- → Cu (s) E o = +0.340 V Cathode b) Al 3+ (aq) 3e - → Al (s) E o = -1.66 V Anode Sn 2+ (aq) → Sn (s) +2e - E o = -0.137 V Cathode 2.a) Ba 2+ (aq) | Ba (s) || Cu (s) | Cu 2+(aq) 2.b) Al (s) | Al 3+ (aq) || Sn 2+ (aq) | Sn (s) 3.a) E o cell = 0.34 - (-2.92) = 3.26 V 3.b) E o cell = -0.137 - (-1.66) = 1.523 V

17. Cell Potential Difference = Voltage If E o cell is positive the reaction is spontaneous and it is a voltaic cell Convert chemical energy into electrical energy If the E o cell is negative, the reaction is non- spontaneous and it is an electrolytic cell. Convert electrical energy into chemical energy