Redox Reactions and Electrochemistry Chapter 19. Applications of Oxidation-Reduction Reactions

Redox Reactions and ElectrochemistryChapter 19

Applications of Oxidation-Reduction Reactions

Batteries

19.6

Leclanché cell

Dry cell

Zn (s) Zn2+ (aq) + 2e-Anode:

Cathode: 2NH4 (aq) + 2MnO2 (s) + 2e- Mn2O3 (s) + 2NH3 (aq) + H2O (l)+

Zn (s) + 2NH4 (aq) + 2MnO2 (s) Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s)

Batteries

Alkaline battery (1.5 V)

Most common nonrechargable battery; provides far superior performance over older “dry cells” that were also based on Zn and MnO2 as the electrochemically active substances

Anode: Zn (s) + 2 OH- (aq) Zn(OH)2 (s) + 2e-

Cathode: 2 MnO2 (s) + 2 H2O (l) + 2e- 2 MnO(OH) (s) + 2OH-(aq)

Batteries

Zn(Hg) + 2OH- (aq) ZnO (s) + H2O (l) + 2e-Anode:

Cathode: HgO (s) + H2O (l) + 2e- Hg (l) + 2OH- (aq)

Zn(Hg) + HgO (s) ZnO (s) + Hg (l)

Mercury Battery

19.6

Used in pacepakers and hearing aids

Batteries

19.6

Anode:

Cathode:

Lead storagebattery

PbO2 (s) + 4H+ (aq) + SO2- (aq) + 2e- PbSO4 (s) + 2H2O (l)4

Pb (s) + SO2- (aq) PbSO4 (s) + 2e-4

Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO2- (aq) 2PbSO4 (s) + 2H2O (l)4

Used in cars and trucks

Batteries

19.6Solid State Lithium Battery

Used in laptops and cell phones

Batteries

19.6

A fuel cell is an electrochemical cell that requires a continuous supply of reactants to keep functioning (not batteries because they are not self-contained systems)

Anode:

Cathode: O2 (g) + 2H2O (l) + 4e- 4OH- (aq)

2H2 (g) + 4OH- (aq) 4H2O (l) + 4e-

2H2 (g) + O2 (g) 2H2O (l)

Used in space vehicles: liquid H2 and O2 are stored as fuel, and the product of the reaction is consumed by the spacecraft crew

Hydrogen Fuel Cells

Corrosion

• Deterioration of metals by an electrochemical process

– Rust on iron– Tarnish on silver– Green patina formed on copper and brass

Corrosion

19.7

• Oxygen gas and water must be present for iron to rust• Reactions are quite complex and not completely undersood, but the

main steps are outlined here• Note that reaction occurs in acidic medium

Corrosion

• In cold climates, salts spread on roadways to melt ice and snow are a major cause of rust formation on automobiles

• Electrical circuit is completed by the migration of electrons and ions

• Therefore, rusting occurs rapidly in salt water

Corrosion PreventionEmploying a sacrificial anode to prevent Fe oxidation

Cathodic Protection of an Iron Storage Tank

19.7

Connecting to a metal that oxidizes more-readily

Corrosion prevention

• A “tin” can is made by applying a thin layer of tin over iron

• Rust is prevented as long as the tin layer remains intact

• However, once the surface has been scratched, rusting occurs rapidly

Coating the surface with a metal that oxidizes less-readily

Electrolysis

• Voltaic cells are based on spontaneous oxidation-reduction reactions

• Conversely, it is possible to use electrical energy to cause nonspontaneous redox reactions to occur

• Such processes, which are driven by an outside source of electrical energy, are called electrolysis reactions and take place in an electrolytic cell

Electricity can be used to decompose molten NaCl into its component elements

Electrolysis

• This is the reason manufacturers of automotive batteries caution against immersing the battery in salt water the standard 12-V car battery has more than enough electromotive force to produce harmful products, such as poisonous Cl2 gas!

Electricity can be used to decompose molten NaCl into its component elements

19.8

A battery (or some other source of direct electrical current) acts as an electron pump, pushing electrons into one electrode and pulling them from the other.

Downs cell Simplified schematic

19.8

The electrodes are inert; they do not undergo a reaction but merely serve as the surface where oxidation and reduction occur.

Downs cell Simplified schematic

Electrolysis

• Note that the cathode of the voltage source is connected to the anode of the electrolytic cell

• And that the anode of the voltage source is connected to the cathode of the electrolytic cell

• Thus the circuit is complete

Electricity can be used to decompose molten NaCl into its component elements

Electrolysis

• Na is not found free in nature due to its great reactivity

• It is obtained commercially by the electrolysis of dry molten sodium chloride

• Sodium is a soft, silvery-white metal which is generally stored in paraffin, as it oxidises rapidly when cut.

Electricity can be used to decompose molten NaCl into its component elements

ElectrolysisElectricity can be used to decompose molten NaCl into its component elements. Why MOLTEN NaCl?

Water undergoes electrolysis

19.8

Water undergoes electrolysis

19.8

Quantitative aspects of electrolysis

19.8

Electroplating uses electrolysis to deposit a thin layer of one metal onto another metal in order to improve beauty or resistance to corrosion.

Quantitative aspects of electrolysis

19.8

An example of electroplating would be depositing a thin layer of nickel onto steel.

• Nickel dissolves from the anode to form Ni2+

(aq)

• At the cathode, Ni2+(aq) is

reduced and forms a nickel “plate” on the cathode

Quantitative aspects of electrolysis

19.8

An example of electroplating would be depositing a thin layer of nickel onto steel.

• These Ernie Ball strings are made from nickel-plated steel wire wrapped around tin plated hex shaped steel core wire. Their nickel-wound sets are by far the most popular, producing a well balanced and all around good sound.

Quantitative aspects of electrolysis

19.8

For any half-reaction, the amount of a substance that is reduced or oxidized in an electrolytic cell is directly proportional to the number of electrons passed into the cell.

• Quantity of charge passing through an electrical circuit, such as that in an electrolytic cell, is generally measured in coulombs

• The charge on 1 mole of electrons is 96,485 C (1 faraday)

• A coulomb is the quantity of charge passing a point in a circuit in 1 s when the current is 1 ampere (A)

• Therefore, number of coulombs passing through a cell can be obtained by multiplying the amperage and the elapsed time in seconds.

Electrolysis and Mass Changes

charge (C) = current (A) x time (s)

1 mole e- = 96,500 C

19.8

How much Ca will be produced in an electrolytic cell of molten CaCl2 if a current of 0.452 A is passed through the cell for 1.5 hours?

Anode:

Cathode: Ca2+ (l) + 2e- Ca (s)

2Cl- (l) Cl2 (g) + 2e-

Ca2+ (l) + 2Cl- (l) Ca (s) + Cl2 (g)

2 mole e- = 1 mole Ca

mol Ca = 0.452Cs

x 1.5 hr x 3600shr 96,500 C

1 mol e-

x2 mol e-

1 mol Cax

= 0.0126 mol Ca

= 0.50 g Ca

19.8