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Galvanic Cells and Electrolytic Cells

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Megan Kaiser

on 2 May 2014

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Transcript of Galvanic Cells and Electrolytic Cells

Galvanic Cells and Electrolytic Cells
Megan Kaiser
AP Chemistry- 1st period
4/29/14
Galvanic, or voltaic cells, produce electricity by using a redox reaction. The redox reaction in a galvanic cell is a spontaneous reaction. For this reason, galvanic cells are commonly used as batteries. Galvanic cell reactions supply energy which is used to perform work. The energy is harnessed by situating the oxidation and reduction reactions in separate containers, joined by an apparatus that allows electrons to flow. A common galvanic cell is the Daniell cell, shown below.
Redox Reactions
Electrochemical reactions involve redox reactions. Redox stands for reduction and oxidation. Reduction is the gain of electrons and involves an decrease in oxidation number. Oxidation is the loss of electrons and an increase in oxidation number. So, in redox reactions electrons are being exchanged as reactants are being converted into products. This electron exchange may be direct, as when copper metal plates out on a piece of zinc, or it may be indirect, as in an electrochemical cell (battery).
Electrolytic Cells
The redox reaction in an electrolytic cell is nonspontaneous. Electrical energy is required to induce the electrolysis reaction. Electrolytic cells use electricity from an external source to produce a desired redox reaction. Electroplating and the recharging of an automobile are examples of electrolytic cells.
An example of an electrolytic cell is shown above, in which molten NaCl is electrolyzed to form liquid sodium and chlorine gas. The sodium ions migrate toward the cathode, where they are reduced to sodium metal. Similarly, chloride ions migrate to the anode and are oxidized to form chlorine gas. This type of cell is used to produce sodium and chlorine. The chlorine gas can be collected surrounding the cell. The sodium metal is less dense than the molten salt and is removed as it floats to the top of the reaction container.
electrodes
: solid portion of the cell that conducts the electrons involved in the redox reaction
electrode

compartments
: solutions in which the electrodes are immersed
salt

bridge
: an inverted U-tube that holds a gel containing a concentrated electrolyte solution, such as KNO3.
anode
: the electrode at which oxidation is taking place
anode compartment
: the electrolyte solution in which the anode is immersed
cathode
: the electrode at which reduction takes place
cathode compartment
: the electrolyte solution in which the cathode is immersed
Key Terms
Cell Potential
The cell potential, Ecell, is the measure of the potential difference between two half cells in an electrochemical cell. The potential difference is caused by the ability of electrons to flow from one half cell to the other. Electrons are able to move between electrodes because the chemical reaction is a redox reaction. This relates to the measurement of the cell potential because the difference between the potential for the reducing agent to become oxidized and the oxidizing agent to become reduced will determine the cell potential. The cell potential (Ecell) is measured in voltage (V), which allows us to give a certain value to the cell potential.
standard cell potential
: E° the potential (voltage) associated with the cell at standard conditions
A few things should be considered when using standard reduction potentials to generate the cell reaction and cell potential:
standard cell potential for a galvanic cell is a positive value, E° > 0
b/c one half-reaction must involve oxidation, one of the half-reactions shown in the table of reduction potentials must be reversed to indicate the oxidation. If the half-reaction is reversed, the sign of the standard reduction potential must be reversed.
b/c oxidation occurs at the anode and reduction at the cathode, the standard cell potential can be calculated from the standard reduction potentials of the two half-reactions involved in the overall reaction by using the equation: E°cell= E°cathode - E°anode > 0
Electrolysis
Electrolysis is one of the most widely used applications of electrolytic cells, and it is the decomposition of a compound. When undergoing electrolysis, several questions may be asked:
1. How long will it take?
2. How much an be produced?
3. What current must be used?
Given any two of these quantities, the third can be calculated, but the balanced half-reaction must be known. Then the following relationships can be applied:
1 Faraday= 96,500 coulombs per mole of electrons
(F= 96,500 C/mol e- OR 96,500 J/V)
1 ampere= 1 coulomb/ second
(A=C/s)
Knowing the amperage and how long it's being applied, the coulombs can be calculated. Then coulombs to moles of electrons. Moles of electrons can be related to moles (and then grams if necessary) of material being electrolyzed through the balanced half-reaction
Nernst Equation
When the cell is not at standard conditions, calculations can not be based on the standard cell potential or standard half-cell potentials. In this case, the actual cell potential (Ecell) can be calculated by the Nernst Equation:

Ecell = E°cell - (RT/nF)lnQ = E°cell - (0.0592/n)logQ (@25°C)

R=ideal gas constant
T= temperature in Kelvin
n= # of electrons transferred
F= Faraday's constant
Q= reaction quotient

Example Problem
Example Problem Solution
Galvanic Cells
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