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Galvanic and Electrolytic Cells

Galvanic and Electrolytic Cells . March 24, 2011. Outline. Last Class : How to balance redox reactions in acidic and basic solutions

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Galvanic and Electrolytic Cells

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  1. Galvanic and Electrolytic Cells March 24, 2011

  2. Outline • Last Class: How to balance redox reactions in acidic and basic solutions • Today’s Lesson: What is a galvanic cell and a electrolytic cell and how do they work? What is the relationship between these cells and the redox reactions we learned about and “balanced” in class? • Tomorrow: How to calculate standard cell potential.

  3. Electrochemistry Concepts • Redox reactions - involve the transfer of electrons from one reactant to another • Electric current is a flow of electrons in a circuit Electrochemistry • Study of the processes involved in converting chemical energy into electrical energy and vice versa

  4. Electrochemical Cells • Redox reactions take place in electrochemical cells. • Many oxidation-reduction reactions occur spontaneously in electrochemical cells • Other oxidation-reduction reactions are not spontaneous in electrochemical cells Galvanic Cells • Spontaneous reactions occur Electrolytic Cells • Non-spontaneous reactions occur

  5. Galvanic Cell (Voltaic Cell) • Converts chemical energy into electrical energy • Key Aspects: • Reactants are not in direct contact with each other • Electrons flow through an external circuit Cu2+(aq) + Zn (s) -------> Cu(s) + Zn2+ (aq) Oxidation: Zn(s) ----> Zn2+(aq) + 2e- Reduction: Cu2+(aq) + 2e- ----> Cu(s)

  6. Cu2+(aq) + Zn (s) ----> Cu(s) + Zn2+ (aq) Oxidation: Zn(s) ----> Zn2+(aq) + 2e- Reduction: Cu2+(aq) + 2e- ----> Cu(s) Galvanic Cell • Two compartments or half-cells, each composed of an electrode dipped in a solution of electrolyte to keep the reaction separate Zinc strip in ZnSO4 solution Copper strip in CuSO4 solution

  7. Galvanic Cell Anode • Oxidation Half-Cell: Anode • Zn(s) ----> Zn2+(aq) + 2e- • Reduction Half-Cell: Cathode • Cu2+(aq) + 2e- ----> Cu(s) How does a redox reaction take place if they aren’t connected? • Connect metal “electrodes” with a conducting wire to allow electrons to flow • However, this eventually leads to a problem. Will eventually get an imbalance of charge • The anode will become more positive (Zn2+ produced) • The cathode will become more negative (Cu2+ removed) • We can solve this imbalance with the use of a salt bridge Zinc strip in zinc sulfate solution Copper strip in copper solution sulfate solution Cathode

  8. External Circuit Galvanic Cell K+ SO42- Salt Bridge – porous barrier that allows the migration of ions in both directions to maintain electrical neutrality • In this example it contains a saturated solution of K2SO4. It doesn’t interfere with the cathode or anode but only maintains electrical charge • Prevents contact between Cu2+ and Zn External Circuit – forces electrons to move from anode to cathode (between 2 electrodes) and powers any electrical device Cu2+ Zn2+ Anode Zn2+ increases SO42- stays the same Positive charge builds up Cathode Cu2+ decreases SO42- stays the same Negative charge builds up

  9. Cell Notation Cu2+(aq) + Zn (s) ----> Cu(s) + Zn2+ (aq) • Allows for a short way to represent a galvanic cell • │ Phase boundary (solid  aqueous) • ││ Porous barrier or salt bridge • Occurs in the following order • Anode │cation of anode ││ cation of cathode │ cathode • Zn │ Zn2+ ││ Cu2+ │ Cu • Question: What will eventually happen to the mass of the Zn and Cu electrodes as this redox reaction proceeds? • Zn anode will decrease in mass • Cu cathode will increase in mass

  10. Summary of Galvanic Cells

  11. The alkaline cell Battery • Type of Galvanic Cell: • Half Reactions : • Oxidation (anode): Zn(s) + 2OH-(aq+) ZnO(s) + H2O(l) +2e- • Reduction (cathode): MnO2(s) + 2(H2O)(l) + 2e-  Mn(OH)2(s) + 2OH(aq) • Overall cell reaction: Zn(s) +MnO2(s) + H2O(l)  ZnO(s) + Mn(OH)2(s)

  12. Electrolytic Cell • Converts electrical energy into chemical energy • Key Aspects: • The process that takes place in an electrolytic cell is called electrolysis. • Overall reaction is non-spontaneous and requires energy for it to occur. • Requirements: • Like a galvanic cell, it requires electrodes, at least one electrolyte, and an external circuit. • Do the half reactions need to be separated?

  13. Electrolytic Cell • Examples: Cu2+(aq) + Zn(s) ---> Cu(s) + Zn2+(aq) -This reaction can be forced to go in reverse if we apply energy Electrolysis of Water • Half Reactions: Oxidation (anode): 2H2O(l)  O2(g) + 4H+(aq) + 4e- Reduction (cathode): 4H2O(l) 4e- 2H2(g) + 4OH- Overall: 2H2O(l)  O2(g) + 2H2(g)

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