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An overview of electrochemical cells, discussing the differences between electrolytic and voltaic cells, electrical conduction, ionic conduction, electrodes, and the processes of redox reactions and standard electrode potentials. It also includes examples of specific electrochemical cells and their reactions.
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Electrolytic Cells: electrical energy from an external source causes a nonspontaneous reaction to occur Voltaic Cells (Galvanic Cells): spontaneous chemical reactions produce electricity and supply it to an external circuit
Electric current represents charge transfer Charges conducted through:
Surfaces upon which oxidation and reduction half reactions occur May or may not participate in the reaction Inert Electrodes – do not participate Ex. Pt, C, Pd Reduction at cathode Oxidation at anode
Electrodes RED CAT And AN OX
Spontaneous oxidation – reduction reactions produce electrical energy Two halves of redox reaction are separated Half cell – contains the oxidized and reduced forms of an element or other complex species
Salt bridge – completes circuit between the two half cells Salt bridge is any medium through which ions can flow Agar + Salt Gelations
Zn metal Cu 2+ ions Zn metal Cu 2+ ions With time, Cu plates onto the Zn metal strip, and Zn strip disappears
http://www.chembio.uoguelph.ca/educmat/chm 105/galvanic/galvanic1.htm http://www.youtube.com/watch?v=0oSqPDD2rM A
Zn – Cu Cu electrode – cathode Cu+2^ is more easily reduced than Zn+ Zn is a stronger reducing agent than Cu Ag – Cu Cu electrode – anode Ag+^ is more easily reduced than Cu+ Cu is a stronger reducing agent than Ag Cathode – Anode are dictated by species present
Strength as oxidizing agents Zn+2^ < Cu+2^ < Ag+ Strength as reducing agents Zn > Cu > Ag
“Every oxidation needs a reduction” e-^ must go somewhere Therefore it is impossible to determine experimentally the potential of a single electrode Establish an arbitrary standard electrode Standard Hydrogen Electrode, SHE
Metal coated with Pt immersed in a 1.0 M H+ solution. H 2 gas is bubbled at 1 atm over the electrode Assigned a potential of 0.000 V **2 H2 H+
(aq, 1 M)(aq, 1 M) + 2e- <----> H+ 2e- <----> H 22 (g, 1 atm)(g, 1 atm) EE° = 0.000V° = 0.000V H H 22 (g, 1 atm(g, 1 atm <----> 2 H<----> 2 H++(aq, 1 M)(aq, 1 M) + 2e-+ 2e- EE° = 0.000V° = 0.000V**
