Chp 17 Electrochemistry.

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Presentation transcript:

Chp 17 Electrochemistry

Galvanic Cells A device in which chemical energy is converted into electrical energy Oxidation – reduction (redox) reactions e- ‘s transfer from reducing agent to oxidizing agent

Galvanic Cell e- e- e- Salt Bridge Why? Zn Cu Zn 2+ Cu 2+ SO4 2- Voltage meter e- e- e- Salt Bridge Why? Zn Cu Zn 2+ SO4 2- Cu 2+ SO4 2- http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf

What is happening to the Zn and the Cu? Zn is making electrons (oxidization) Cu is gaining electrons (reduction) Anode where electrons are produced. Cathode where electrons are used. Cell potential cell (volt) or (V) Zn disappearing Forming Cu

Standard Reduction Potentials Anode: Zn(s)  Zn 2+(aq) + 2 e- Cathode: Cu 2+(aq) + 2 e-  Cu (s) ocell = ocat + oan ocell = (+0.76V) + (+0.34V) = +1.10 V Must be positive

Examples Fe 3+(aq) + Cu (s)  Cu 2+(aq) + Fe 2+(aq) Half reactions Red: ( Fe 3+ + 1e -  Fe 2+) Ox: Cu  Cu 2+ + 2e – 2Fe 3+(aq) + Cu (s)  Cu 2+(aq) + 2Fe 2+(aq) ocell = ocat + oan ocell = 0.43 V o = 0.77 V 2 o = - 0.34 V Opposite sign because of Ox = 0.77 V + (-0.34 V)

Line Notation Zn (s) + Cu 2+(aq)  Zn 2+(aq) + Cu (s) Zn(s) l Zn 2+(aq) ll Cu 2+(aq) l Cu(s) electrode electrode