Electrochemistry Lesson 2

Strength of Metals in Redox Look at the standard reduction table (next slide). For any two metals in the table, the metal with the larger value is reduced while the metal with the smaller value is oxidized. (the table is a measurement of “reduction potential” the higher the number the more potential for reduction) Which is more readily oxidized, magnesium or lead?

Strength of Metals in Redox Look at the standard reduction table (next slide). For any two metals in the table, the metal with the larger value is reduced while the metal with the smaller value is oxidized. (the table is a measurement of “reduction potential” the higher the number the more potential for reduction) Which is more readily oxidized, magnesium or lead? Magnesium

Voltaic/Galvanic Cells Electrochemical cells used to convert chemical energy into electrical energy. The electrical energy is produced by spontaneous redox reactions in the cell. A voltaic cell is made up of two half cells connected by a salt bridge that ions travel through. Each half cell contains either an anode or a cathode, connected by a wire that electrons flow through.

Oxidation Reduction

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Electrons flow from the anode to the cathode ReDox The reduction half-reaction takes place at the cathode (+) and the oxidation half-reaction takes place at the anode (-). RedCat AnOx Electrons flow from the anode to the cathode

Standard Cell Potential A voltaic cell’s ability to produce an electric current. E°cell = E°red + E°ox E°cell is the standard reduction potential of the voltaic cell measured in volts (V). E°red is the standard reduction potential of the reduction half-reaction. E°ox is the standard reduction potential of the oxidation half-reaction. (This value will have the opposite sign from the reduction potential table.) E°cell must be positive for the reaction to be spontaneous.

Calculating Cell Potential Determine the voltage produced by the following redox reaction. Zn(s) + 2 H+(aq)  Zn2+(aq) + H2(g) Oxidation: Zn(s)  Zn2+(aq) + 2 e- Reduction: 2 H+(aq) + 2 e-  H2(g) E°cell = E°red + E°ox E°cell = 0.00 V + (0.76 V) = +0.76 V

Often you will be asked to determine the standard cell potential given the two half- cell reactions. First, determine which reaction is the reduction and which is the oxidation. The half-cell with the larger reduction potential is the one at which reduction takes place. Calculate the standard cell potential using the equation and the values from the standard reduction table. Note: The standard reduction potential of a half-reaction is never multiplied by a coefficient or changed in sign.

Sample Problem Li+(aq) + e-  Li(s) Mg2+(aq) + 2 e-  Mg(s) A voltaic cell is constructed using the following half reactions. Determine the standard cell potential. Li+(aq) + e-  Li(s) Mg2+(aq) + 2 e-  Mg(s) Ox: Li(s)  Li+(aq) + e- Red: Mg2+(aq) + 2 e-  Mg(s) E°cell = -2.37 V + (3.05 V) = +0.68 V

Sample Problem Determine whether the following redox reaction is spontaneous. 2 Ag(s) + Zn2+(aq)  2 Ag+(aq) + Zn(s) Ox: 2 Ag(s)  2 Ag+(aq) + 2 e- Red: Zn2+(aq) + 2 e-  Zn(s) E°cell = -0.76 V + (- 0.80 V) = -1.56 V Not Spontaneous

Batteries A battery is any electrochemical cell that spontaneously produces electrical energy. A battery has stored chemical energy that can be converted to electrical energy by a chemical reaction. A battery is made up of one or more voltaic cells. An automobile battery is made up of six cells. A battery without an electrolyte solution is a dry cell.

Batteries

Alkaline Batteries

Hydrogen Fuel Cells

Electrolytic Cells Electrolytic cells are electrochemical cells that do not operate spontaneously. The process is referred to as electrolysis. A source of electricity is required to drive an electrolytic cell. An example of an electrolysis reaction is the recharging of the battery in a cell phone.