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Chem 1062 Notes ch 21: Electrochemistry

by: mandygh926

Chem 1062 Notes ch 21: Electrochemistry CHEM 1062

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These notes include the topics: Cell Potential, Anodes and Cathodes, Nernst Equation, Voltaic and Electrolytic Cells, Half Cells, and Balancing Redox Reactions.
Chemical Principles II
Doreen Leopold
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This 2 page Bundle was uploaded by mandygh926 on Sunday July 31, 2016. The Bundle belongs to CHEM 1062 at University of Minnesota taught by Doreen Leopold in Spring 2016. Since its upload, it has received 7 views. For similar materials see Chemical Principles II in Chemistry at University of Minnesota.

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Date Created: 07/31/16
Chemistry Notes Ch. 21: Electrochemistry: Chemical Change and Electrical Work Includes: Cell Potential, Anodes and Cathodes, Nernst Equation, Voltaic and Electrolytic Cells, Half Cells, Balancing Redox Reactions, and Spontaneity. Corresponding Textbook: Chemistry; The Molecular Nature of Matter and Change Key Terms:  Cell potential: (Cell) depends on the composition of cells, concentration, and temp. o E°cell=0 only if both half-cells are made of the same material. o Reverse half-rxns do not reverse the sign of E° half-cell o A metal can reduce another species if E°cell is positive for the overall reaction  Standard electrode potential: species lower on the table will reduce species higher on the table  Nernst Equation: calculates the cell potential when given the standard cell potential and initial conditions for the concentrations. Given by the following equation, where Q is the initial condition concentrations, and n is the number of moles of electrons transferred: −0.0592V Ecell=E °cell= n logQ  Redox Reactions: Occur when there is a species being reduced and another species being oxidized.  Voltaic cell (galvanic cell): uses a spontaneous reaction to generate electrical energy. It does work on the surroundings.  Electrolytic cell: uses electrical energy to drive a nonspontaneous reaction to generate electrical energy.  Anode: where ox. Half-rxn occurs. e- leave ox half-cell at the anode  Cathode: where the red. Half-rxn occurs. e- enter red. Half-cell at cathode.  Half cell: consists of 1 electrode dipped into an electrolyte solution. It is used in any voltaic cell to separate half-reactions into different containers. Key Concepts:  Spontaneous Cell Potential: Ecell > 0 for spontaneous processes; e- flow from negative to positive (anode to cathode) to higher concentration. E°cell=E°cathode ¿.)−E°anode (ox.=E°unknown−E°reference  Increasing cell Potential: A voltaic cell always runs spontaneously towards producing + Ecell. Therefore, to increase cell potential, increase the amount (concentration) of reactants in the cathode.  Comparing Cell Potentials: For Q= the initial cell condition concentrations, o If Q=1 then Ecell=E°cell o If Q<1 then Ecell> E°cell  Anagram to help with redox reactions: OIL RIG = oxidation loses e- and reduction gains e-  Half-reaction method for balancing redox reactions: 1. Divide into half rxns 2. Balance atoms and charges (all others before O Then H), balance charges w e- on left side of redox rxn and right side of ox. Rxn 3. If necessary: multiply 1 or both rxns by integer so e- gained= e- lost 4. Add together balanced half reactions 5. For Acidic solutions: use H2O to balance the O’s 6. Add H+ to balance H 7. Add e- to balance the charges **Note: For Basic solutions, balance them like acids, but then add an OH- to both sides of the equation for every H+


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