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electrochemistry
the branch of chemistry that deals with the interconversion between electrical energy and chemical energy
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Redox reactions
reactions involving the trasfer of electrons from one substance to another
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Steps for balancing redox reactions
- 1. Write the unbalanced ionic equation
- 2. Separate into 2 half-reactions
- 3. Balance atoms other than H & O
- 4. Add H2O to balance O atoms
- 5. Add H+ to balance H (if basic add OH to balance H+)
- 6. Add e- to 1 side of each 1/2 rxn to balance changes
- 7. Equalize the # of e- in the 2 1/2 rxns by multiplying by proper coefficients
- 8. Add rxns together, e- must cancel
- 9. Verify atoms and sum of charges are balanced
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galvanic cell
aka voltaic cell - the experimental apparatus for generating electricity through the use of spontaneous redox reaction
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salt bridge
contains an inert electrolytic solution that allows the charges to stay balanced
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anode
- where oxidation takes place, usually listed 1st
- electrons flow from here to cathode
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cathode
- where reduction takes place, usually listed 2st
- electrons flow from anode to here
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E
- Electromotive force
- Cell voltage
- cell potential
- emf
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SHE
Hydrogen Electrode Operating under Standard-State Conditions
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standard reduction potential E*
the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm
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standard oxidation potential
the voltage associated with an oxidation reaction at an electrode when all solutes are 1 M and all gases are at 1 atm (- of the standard reduction potential)
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standard cell potential E*cell
- standard emf - the sum of the standard oxidation potential and standard reduction potential
- E*cell = E*ox + E* red
- MAKE SURE of order (solid on right is reduction)
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equation to relate free-energy change to the emf of a cell
^ G = -n F Ecell
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equation relating standard emf of the cell to the equilibrium constant
E*cell = 0.0257V In K
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Values of ^G*, K, E*cell for a spontaneous reactions
- ^G* < 0 (-)
- K > 1 (big)
- E*cell > 0 (+)
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Values of ^G*, K,
E*cell for a non-spontaneous reactions
- ^G* > 0 (+)
- K < 1 (small)
- E*cell< 0 (-)
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Values of ^G*, K,
E*cell for a reaction at equilibrium
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Nernst equation
relates emf to concentrations and non-standard conditions
E = E* - RT/ nF In Q
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dry cell battery
Zinc, Manganese dioxide, ammonium chloride
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Mercury battery
expensive - used in medicine and electronic industries
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Lead Storage battery
automobiles, 6 cells
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fuel cell
a galvanic cell that requires a continuous supply of reactants to keep functioning
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corrosion
the deterioration of metals by an electrochemical process
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electrolysis
the process in which electrical energy is used to cause a nonspontaneous chemical reaction to occur
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Amperes
(A) 1C = 1A x 1sec
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faraday
- the charge on a mole of electrons, it is synonymous with 1 mole of e-
- Coulombs x Volts = Joules
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