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Electrochemical Cells (voltaic cells)

Electrochemical Cells (voltaic cells). Electrochemical Cells :. **Spontaneous Redox. Zn 0 + Cu +2  Zn +2 + Cu 0. Table J (Activity Series) :. Single Replacement  the more active metal replaces the less active metal. Oxidation: Zn  Zn 2+ + 2e -. Reduction: Cu +2 + 2e -  Cu 0.

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Electrochemical Cells (voltaic cells)

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  1. Electrochemical Cells (voltaic cells)

  2. Electrochemical Cells: **Spontaneous Redox Zn0 + Cu+2 Zn+2 + Cu0

  3. Table J (Activity Series): Single Replacement  the more active metal replaces the less active metal Oxidation: Zn Zn2+ + 2e- Reduction: Cu+2 + 2e- Cu0

  4. Voltaic Cells: **When the half reactions are separated and they force a flow of electricity. *Salt bridge allows ions to move between half cells.

  5. Anode: Where oxidation takes place Cathode: Where reduction takes place Voltaic Cell Animation

  6. Summary • Electrons are produced at the Zn electrode (anode) • Zn Zn2+ + 2e- Electrons produced causes negative charge • 2. Electrons leave Zn electrode and pass through external circuit (wire)

  7. Summary • 3. Electrons at Cu electrode are used to reduce (cathode) • Cu+2 + 2e- Cu0 Electrons used up causes positive charge • 4. Salt bridge completes the circuit allowing ions to move

  8. Dry Cells Zinc-Carbon batteries Anode (-): Oxidation Zn  Zn2+ + 2e- Cathode (+): Reduction Mn4++e- Mn3+

  9. Fuel Cells Oxidation of a fuel to produce electricity 2H2 + O2 2H2O Anode (-): Oxidation 2H20 4H+ + 4e- Cathode (+): Reduction O2 + 4e-  2O2-

  10. Electrolytic Cells **Requires electricity for the reaction (called electrolysis) **Reverse of electrochemical cell 2NaCl  2Na + Cl2 Used to get active metals: Na, K, Mg Cathode (-): Reduction Na+(aq)+ e-  Na(s) Anode (+): Oxidation 2Cl-(aq) Cl2(g)+ 2e-

  11. Electrolysis of Water

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