Transcript document

Redox Reactions and
Electrochemistry
Chapter 19
Applications of
Oxidation-Reduction
Reactions
Batteries
Dry cell
Leclanché cell
Anode:
Cathode:
Zn (s)
2NH+4 (aq) + 2MnO2 (s) + 2e-
Zn (s) + 2NH4 (aq) + 2MnO2 (s)
Zn2+ (aq) + 2eMn2O3 (s) + 2NH3 (aq) + H2O (l)
Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s)
19.6
Batteries
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
Alkaline battery (1.5 V)
Anode:
Cathode:
Zn (s) + 2 OH- (aq)
Zn(OH)2 (s) + 2e-
2 MnO2 (s) + 2 H2O (l) + 2e-
2 MnO(OH) (s) + 2OH-(aq)
Batteries
Used in pacepakers and hearing aids
Mercury Battery
Anode:
Cathode:
Zn(Hg) + 2OH- (aq)
HgO (s) + H2O (l) + 2eZn(Hg) + HgO (s)
ZnO (s) + H2O (l) + 2eHg (l) + 2OH- (aq)
ZnO (s) + Hg (l)
19.6
Batteries
Used in cars and trucks
Lead storage
battery
Anode:
Cathode:
Pb (s) + SO2-4 (aq)
PbSO4 (s) + 2e-
PbO2 (s) + 4H+ (aq) + SO24 (aq) + 2e
Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO2(aq)
4
PbSO4 (s) + 2H2O (l)
2PbSO4 (s) + 2H2O (l)
19.6
Batteries
Used in laptops and cell phones
Solid State Lithium Battery
19.6
Batteries
Used in space vehicles: liquid H2 and O2
are stored as fuel, and the product of the
reaction is consumed by the spacecraft
crew
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:
2H2 (g) + 4OH- (aq)
O2 (g) + 2H2O (l) + 4e2H2 (g) + O2 (g)
4H2O (l) + 4e4OH- (aq)
2H2O (l)
19.6
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
• 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
19.7
• Note that reaction occurs in acidic medium
Corrosion
• Electrical circuit is completed
by the migration of electrons
and ions
• Therefore, rusting occurs
rapidly in salt water
• In cold climates, salts
spread on roadways to
melt ice and snow are a
major cause of rust
formation on automobiles
Corrosion Prevention
Employing a sacrificial anode to prevent Fe oxidation
Cathodic Protection of an Iron Storage Tank
Connecting to a metal that oxidizes more-readily
19.7
Corrosion prevention
Coating the surface with a metal that oxidizes less-readily
• 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
Electrolysis
Electricity can be used to decompose molten NaCl into its
component elements
•
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
Electrolysis
Electricity can be used to decompose molten NaCl into its
component elements
• 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!
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
19.8
Electrolysis
Electricity can be used to decompose molten NaCl into its
component elements
• 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
Electrolysis
Electricity can be used to decompose molten NaCl into its
component elements
• 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, silverywhite metal which is
generally stored in
paraffin, as it oxidises
rapidly when cut.
Electrolysis
Electricity 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
Electroplating uses electrolysis to deposit a thin layer of one
metal onto another metal in order to improve beauty or
resistance to corrosion.
19.8
Quantitative aspects of electrolysis
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
19.8
Quantitative aspects of electrolysis
An example of electroplating would be depositing a thin layer of
nickel onto steel.
• These Ernie Ball strings
are made from nickelplated steel wire
wrapped around tin
plated hex shaped steel
core wire. Their nickelwound sets are by far
the most popular,
producing a well
balanced and all around
good sound.
19.8
Quantitative aspects of electrolysis
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.
19.8
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:
2Cl- (l)
Ca2+ (l) + 2eCa2+ (l) + 2Cl- (l)
Cl2 (g) + 2eCa (s)
Ca (s) + Cl2 (g)
2 mole e- = 1 mole Ca
C
s 1 mol e- 1 mol Ca
mol Ca = 0.452
x 1.5 hr x 3600 x
x
s
hr 96,500 C 2 mol e= 0.0126 mol Ca
= 0.50 g Ca
19.8