1 / 11

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.

amish
Télécharger la présentation

Electrochemical Cells (voltaic cells)

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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

More Related