Roy Kennedy

1
Introductory Chemistry, 3rd EditionNivaldo Tro
Chapter 16 Oxidation and Reduction

• Roy Kennedy
• Massachusetts Bay Community College
• Wellesley Hills, MA

2009, Prentice Hall
2
OxidationReduction Reactions

• Oxidationreduction reactions are also called
  redox reactions.
• All redox reactions involve the transfer of
  electrons from one atom to another.
• Spontaneous redox reactions are generally
  exothermic, and we can use their released energy
  as a source of energy for other applications.
• Convert the heat of combustion into mechanical
  energy to move our cars.
• Use electrical energy in a car battery to start
  our car engine.

3
Combustion Reactions

• Combustion reactions are always exothermic.
• In combustion reactions, O2 combines with all the
  elements in another reactant to make the
  products.
• 4 Fe(s) 3 O2(g) ? 2 Fe2O3(s) energy
• CH4(g) 2 O2(g) ? CO2(g) 2 H2O(g) energy

4
Reverse of Combustion Reactions

• Since combustion reactions are exothermic, their
  reverse reactions are endothermic.
• The reverse of a combustion reaction involves the
  production of O2.
• energy 2 Fe2O3(s) ? 4 Fe(s) 3 O2(g)
• energy CO2(g) 2 H2O(g) ? CH4(g) 2 O2(g)
• Reactions in which O2 is gained or lost are redox
  reactions.

5
Oxidation and ReductionOne Definition

• When an element attaches to an oxygen during the
  course of a reaction it is generally being
  oxidized.
• In CH4(g) 2 O2(g) ? CO2(g) 2 H2O(g), C is
  being oxidized in this reaction, but H is not.
• When an element loses an attachment to oxygen
  during the course of a reaction, it is generally
  being reduced.
• In 2 Fe2O3(s) ? 4 Fe(s) 3 O2(g), the Fe is
  being reduced.
• One definition of redox is the gain or loss of O,
  but it is not the best.

6
Another OxidationReduction

• Consider the following reactions
• 4 Na(s) O2(g) ? 2 Na2O(s)
• 2 Na(s) Cl2(g) ? 2 NaCl(s)
• The reaction involves a metal reacting with a
  nonmetal.
• In addition, both reactions involve the
  conversion of free elements into ions.
• 4 Na(s) O2(g) ? 2 Na2O(s)
• 2 Na(s) Cl2(g) ? 2 NaCl(s)

7
Oxidation and ReductionAnother Definition

• In order to convert a free element into an ion,
  the atoms must gain or lose electrons.
• Of course, if one atom loses electrons, another
  must accept them.
• Reactions where electrons are transferred from
  one atom to another are redox reactions.
• Atoms that lose electrons are being oxidized,
  atoms that gain electrons are being reduced.

2 Na(s) Cl2(g) ? 2 NaCl(s) Na ? Na 1 e
(oxidation) Cl2 2 e ? 2 Cl (reduction)
8
PracticeIdentify the Element Being Oxidized and
the Element Being Reduced.

• 2 C(s) O2(g) ? 2 CO(g)
• Mg(s) Cl2(g) ? MgCl2(s)
• Mg(s) Fe2(aq) ? Mg2(aq) Fe(s)

9
PracticeIdentify the Element Being Oxidized and
the Element Being Reduced, Continued.

• 2 C(s) O2(g) ? 2 CO(g)
• Mg(s) Cl2(g) ? MgCl2(s)
• Mg(s) Fe2(aq) ? Mg2(aq) Fe(s)

C is oxidized because it is gaining an attachment
to O. O is reduced there has to be reduction
and its the only other element.
2
-
0
0
Mg is oxidized because it is becoming a cation by
losing electrons. Cl is reduced because it is
becoming an anion by gaining electrons.
Mg is oxidized because it is becoming a cation by
losing electrons. Fe2 is reduced because it is
gaining electrons to become neutral.
10
OxidationReduction

• Oxidation and reduction must occur
  simultaneously.
• If an atom loses electrons, another atom must
  take them.
• The reactant that reduces an element in another
  reactant is called the reducing agent.
• The reducing agent contains the element that is
  oxidized.
• The reactant that oxidizes an element in another
  reactant is called the oxidizing agent.
• The oxidizing agent contains the element that is
  reduced.

2 Na(s) Cl2(g) ? 2 NaCl(s) Na is oxidized, Cl
is reduced. Na is the reducing agent, Cl2 is the
oxidizing agent.
11
PracticeIdentify the Oxidizing and Reducing
Agents.

• 2 C(s) O2(g) ? 2 CO(g)
• Mg(s) Cl2(g) ? MgCl2(s)
• Mg(s) Fe2(aq) ? Mg2(aq) Fe(s)

C is oxidized because it is gaining attachment to
O. O is reduced there has to be reduction and
its the only other element.
2
-
0
0
Mg is oxidized because it is becoming a cation by
losing electrons. Cl is reduced because it is
becoming an anion by gaining electrons.
Mg is oxidized because it is becoming a cation by
losing electrons. Fe2 is reduced because it is
gaining electrons to become neutral.
12
PracticeIdentify the Oxidizing and Reducing
Agents, Continued.

• 2 C(s) O2(g) ? 2 CO(g)
• Mg(s) Cl2(g) ? MgCl2(s)
• Mg(s) Fe2(aq) ? Mg2(aq) Fe(s)

C is the reducing agent because it contains the
element that is oxidized. O is the oxidizing
agent because it contains the element that is
reduced.
2
-
0
0
Mg is the reducing agent because it contains the
element that is oxidized. Cl2 is the oxidizing
agent because it contains the element that is
reduced.
Mg is the reducing agent because it contains the
element that is oxidized. Fe2 is the oxidizing
agent because it contains the element that is
reduced.
13
Electron Bookkeeping

• For reactions that are not metal nonmetal, or
  do not involve O2, we need a method for
  determining how the electrons are transferred.
• Chemists assign a number to each element in a
  reaction called an oxidation state that allows
  them to determine the electron flow in the
  reaction.
• Although they look like them, oxidation states
  are not ion charges!
• Oxidation states are imaginary charges assigned
  based on a set of rules.
• Ion charges are real, measurable charges.

14
Rules for Assigning Oxidation States

• Rules are in order of priority.
• Free elements have an oxidation state 0.
• Na(s) 0 and Cl2(g) 0 in 2 Na(s) Cl2(g)?
• 2 NaCl(s).
• Monoatomic ions have an oxidation state equal to
  their charge.
• Na 1 and Cl -1 in NaCl(s).
• a. The sum of the oxidation states of all the
  atoms or ions in a compound is 0.
• Na 1 and Cl -1 in NaCl, and (1) (-1) 0.

15
Rules for Assigning Oxidation States, Continued

• b. The sum of the oxidation states of all the
  atoms in a polyatomic ion equals the charge on
  the ion.
• N 5 and O -2 in NO3, (5) 3(-2) -1.
• a. Group I metals have an oxidation state of 1
  in all their compounds.
• Na 1 in NaCl.
• b. Group II metals have an oxidation state of
  2 in all their compounds.
• Mg 2 in MgCl2.

16
Rules for Assigning Oxidation States, Continued

• In their compounds, nonmetals have oxidation
  states according to the table below.
• Nonmetals higher on the table take priority.

17
PracticeAssign an Oxidation State to Each
Element in the Following

• F2
• Mg2
• KCl
• SO2
• PO43-
• BaO2

18
PracticeAssign an Oxidation State to Each
Element in the Following, Continued

• F2 F 0 (Rule 1)
• Mg2 Mg 2 (Rule 2)
• KCl K 1 (Rule 4a) and Cl -1 (Rule 5)
• SO2 O -2 (Rule 5) and S 4 (Rule 3a)
• PO43- O -2 (Rule 5) and P 5 (Rule 3b)
• BaO Ba 2 (Rule 4b) and O -2

19
Oxidation and ReductionA Better Definition

• Oxidation occurs when an atoms oxidation state
  increases during a reaction.
• Reduction occurs when an atoms oxidation state
  decreases during a reaction.

CH4 2 O2 ? CO2 2 H2O -4 1
0 4 2 1 -2
20
PracticeAssign Oxidation States and Identify the
Oxidizing and Reducing Agents in Each of the
Following

• 3 H2S 2 NO3 2 H 3 S 2 NO 4 H2O
• MnO2 4 HBr MnBr2 Br2 2 H2O

21
PracticeAssign Oxidation States and Identify the
Oxidizing and Reducing Agents in Each of the
Following, Continued
reducing agent
oxidizing agent

• 3 H2S 2 NO3 2 H 3 S 2 NO 4 H2O
• MnO2 4 HBr MnBr2 Br2 2 H2O

1 -2 5 -2 1 0
2 -2 1 -2
Oxidizing agent
reducing agent
4 -2 1 -1 2 -1 0
1 -2
22
Balancing Redox Reactions

• Some redox reactions can be balanced by the
  method we previously used, but many are hard to
  balance using that method.
• Many are written as net ionic equations.
• Many have elements in multiple compounds.
• The main principle is that electrons are
  transferred, so if we can find a method to keep
  track of the electrons, it will allow us to
  balance the equation.

23
Balancing Redox Reactions by the Half-Reaction
Method

• In this method, the reaction is broken down into
  two half-reactions, one for oxidation and another
  for reduction.
• Each half-reaction includes electrons.
• Electrons go on the product side of the oxidation
  half-reactionloss of electrons.
• Electrons go on the reactant side of the
  reduction half-reactiongain of electrons.
• Each half-reaction is balanced for its atoms.
• Then the two half-reactions are adjusted so that
  the electrons lost and gained will be equal when
  added.

24
Balancing Redox Reactions in Acidic Solution

• Assign oxidation states and determine element
  oxidized and element reduced.
• Separate into oxidation and reduction
  half-reactions.

25
Balancing Redox Reactions in Acidic Solution,
Continued

• 3. Balance half-reactions by mass.
• First balance atoms other than O and H.
• Balance O by adding H2O to side that lacks O.
• Balance H by adding H to side that lacks H.
• Finished if in acidic solution.
• If in basic solution, add enough OH- to
  neutralize the H, rewrite H OH- as H2O.
• Add to both sides.
• Then cancel H2O on both sides.

Fe2 ? Fe3
Tro's Introductory Chemistry, Chapter 16
26
Balancing Redox Reactions in Acidic Solution,
Continued

• 4. Balance each half-reaction, with respect to
  charge, by adjusting the numbers of electrons.
• Electrons on product side for oxidation.
• Electrons on reactant side for reduction.
• 5. Balance electrons between half-reactions.
• 6. Add half-reactions, canceling electrons and
  common species.
• 7. Check.

Fe2 ? Fe3 1 e-
MnO4 8H 5 e- ? Mn2 4H2O
x 5
5 Fe2 MnO4 8H ? Mn2 4H2O 5 Fe3
27
PracticeBalance the Following EquationCu I2
? Cu2 I
28
PracticeBalance the Following Equation,
ContinuedCu I2 ? Cu2 I
1
0
2
-1
oxid
red
oxid Cu ? Cu2
red I2 ? I
oxid Cu ? Cu2
red I2 ? 2 I
oxid Cu ? Cu2 1 e-
red I2 2 e- ? 2 I
oxid Cu ? Cu2 1 e- x 2
red I2 2 e- ? 2 I
2 Cu I2 ? 2 Cu2 I2
29
PracticeBalance the Following Equationin Acidic
Solution I Cr2O72- ? Cr3 I2
30
PracticeBalancing Redox Reactions

• Assign oxidation states and determine element
  oxidized and element reduced.
• Separate into oxidation and reduction
  half-reactions.

oxid I- ? I2
red Cr2O72 ? Cr3
Tro's Introductory Chemistry, Chapter 16
30
31
PracticeBalancing Redox Reactions, Continued

• 3. Balance half-reactions by mass.
• First balance atoms other than O and H.
• Balance O by adding H2O to side that lacks O.
• Balance H by adding H to side that lacks H.
• Finished if in acidic solution.
• If in basic solution, add enough OH- to
  neutralize the H, rewrite H OH- as H2O.
• Add to both sides.
• Then cancel H2O on both sides.

oxid I- ? I2
oxid 2 I- ? I2
red Cr2O72 ? Cr3
red Cr2O72 ? 2 Cr3
red Cr2O72 ? 2Cr3 7H2O
Cr2O72 14H ? 2Cr3 7H2O
Tro's Introductory Chemistry, Chapter 16
31
32
PracticeBalancing Redox Reactions, Continued

• 4. Balance each half-reaction, with respect to
  charge, by adjusting the numbers of electrons.
• Electrons on product side for oxidation.
• Electrons on reactant side for reduction.
• 5. Balance electrons between half-reactions.
• 6. Add half-reactions, canceling electrons and
  common species.
• 7. Check.

2 I- ? I2
2 I- ? I2 2e-
Cr2O72 14H ? 2Cr3 7H2O
Cr2O72 14H 6e-? 2Cr3 7H2O
x 3
2 I- ? I2 2e-
6 I- ? 3 I2 6e-
Cr2O72 14H 6e-? 2Cr3 7H2O
Cr2O72 14H 6 I-? 2Cr3 7H2O 3 I2
Tro's Introductory Chemistry, Chapter 16
32
33
Balancing Redox Reactions in Basic Solution

• Assign oxidation states and determine element
  oxidized and element reduced.
• Separate into oxidation and reduction
  half-reactions.

34
Balancing Redox Reactions in Basic Solution,
Continued
CN- ? CNO- CN- H2O ? CNO- CN- H2O ? CNO- 2
H

• 3. Balance half-reactions by mass.
• First balance atoms other than O and H.
• Balance O by adding H2O to side that lacks O.
• Balance H by adding H to side that lacks H.
• Finished if in acidic solution.
• If in basic solution, add enough OH- to
  neutralize the H, rewrite H OH- as H2O.
• Add to both sides.
• Then cancel H2O on both sides.

CN- H2O 2OH- ? CNO- 2H2O
CN- H2O 2OH- ? CNO- 2H 2OH-
CN- 2OH- ? CNO- H2O
MnO4 ? MnO2
MnO4 ? MnO2 2H2O
MnO4 4H ? MnO2 2H2O
MnO4 4H 4OH- ? MnO2 2H2O 8OH-
MnO4 4H2O ? MnO2 2H2O 8OH-
MnO4 2H2O ? MnO2 8OH-
35
Balancing Redox Reactions in Basic Solution,
Continued

• 4. Balance each half-reaction with respect to
  charge by adjusting the numbers of electrons.
• Electrons on product side for oxidation.
• Electrons on reactant side for reduction.
• 5. Balance electrons between half-reactions.
• 6. Add half-reactions, canceling electrons and
  common species.
• 7. Check.

CN- 2OH- ? CNO- H2O
CN- 2OH- ? CNO- H2O 2e-
MnO4 2H2O ? MnO2 4OH-
MnO4 2H2O 3e- ? MnO2 4OH-
x 3
CN- 2OH- ? CNO- H2O 2e-
MnO4 2H2O 3e- ? MnO2 4OH-
x 2
3CN- 6OH- ? 3CNO- 3H2O 6e-
2MnO4 4H2O 6e- ? 2MnO2 8OH-
3CN 2MnO4 H2O ? 3CNO 2MnO2 2OH
36
PracticeBalance the Following Equationin Basic
Solution I Cr2O72- ? Cr3 I2
Tro's Introductory Chemistry, Chapter 16
36
37
Balancing Redox Reactions

• Assign oxidation states and determine element
  oxidized and element reduced.
• Separate into oxidation and reduction
  half-reactions.

oxid I- ? I2
red Cr2O72 ? Cr3
Tro's Introductory Chemistry, Chapter 16
37
38
Balancing Redox Reactions,Continued

• 3. Balance half-reactions by mass.
• First balance atoms other than O and H.
• Balance O by adding H2O to side that lacks O.
• Balance H by adding H to side that lacks H.
• Finished if in acidic solution.
• If in basic solution, add enough OH- to
  neutralize the H, rewrite H OH- as H2O.
• Add to both sides.
• Then cancel H2O on both sides.

oxid I- ? I2
oxid 2 I- ? I2
red Cr2O72 ? Cr3
red Cr2O72 ? 2 Cr3
red Cr2O72 ? 2Cr3 7H2O
Cr2O72 14H ? 2Cr3 7H2O
Cr2O72 14H2O ? 2Cr3 7H2O 14OH-
Cr2O72 14H 14OH- ? 2Cr3 7H2O 14OH-
Cr2O72 7H2O ? 2Cr3 14OH-
38
39
Balancing Redox Reactions,Continued

• 4. Balance each half-reaction with respect to
  charge by adjusting the numbers of electrons.
• Electrons on product side for oxidation.
• Electrons on reactant side for reduction.
• 5. Balance electrons between half-reactions .
• 6. Add half-reactions, canceling electrons and
  common species.
• 7. Check.

2 I- ? I2
2 I- ? I2 2e-
Cr2O72 7H2O ? 2Cr3 14OH-
Cr2O72 7H2O 6e-? 2Cr3 14OH-
x 3
2 I- ? I2 2e-
6 I- ? 3 I2 6e-
Cr2O72 7H2O 6e-? 2Cr3 14OH-
Cr2O72 7H2O 6 I-? 2Cr3 14OH- 3 I2
Tro's Introductory Chemistry, Chapter 16
39
40
Will a Reaction Take Place?

• Reactions that are energetically favorable are
  said to be spontaneous.
• They can happen, but the activation energy may be
  so large that the rate is very slow.
• The relative reactivity of metals can be used to
  determine if some redox reactions are
  spontaneous.

41
Single Displacement Reactions

• Also known as single replacement reactions.
• A more active free element displaces a less
  active element in a compound.
• Metals displace metals or H.
• Cu 2 AgNO3 Cu(NO3)2 2 Ag
• Mg 2 HCl MgCl2 H2
• Nonmetals displace nonmetals.
• 2 KI Br2 2 KBr I2
• Carbon displaces metals from oxides.
• 3 C Fe2O3 3 CO 2 Fe
• Always redox.

42
Tendency to Lose Electrons

• Some metals have a greater tendency to lose
  electrons than others.
• Metallic-free elements are always oxidized.
• The greater the tendency of a metal to lose
  electrons, the easier it is to oxidize.
• The greater the tendency of a metal to lose
  electrons, the harder it is to reduce its
  cations.
• If Metal A has a greater tendency to lose
  electrons than Metal B, then
• A(s) B(aq) ? A(aq) B(s),
• but A(aq) B(s) ? no reaction.

Tro's Introductory Chemistry, Chapter 16
42
43
Activity Series of Metals

• Listing of metals by reactivity.
• Free metal higher on the list displaces metal
  cation lower on the list.
• Metals above H will dissolve in acid

Zn Fe2 Fe Zn2
Cu Fe2 no reaction
44
Mg is above Cu on the activity series.
45
Table of Oxidation Half-Reactions
45
46
Table of Oxidation Half-Reactions, Continued

• Any oxidation half-reaction that is higher on the
  list will give a spontaneous reaction when
  combined with the reverse of a half-reaction that
  is lower on the list.
• The reverse of an oxidation half-reaction is a
  reduction half-reaction.
• Metals will dissolve in acid if their oxidation
  half-reaction is above H2 ? 2H 2e-.

Tro's Introductory Chemistry, Chapter 16
46
47
ExampleComplete and Balance the Following
ReactionAl(s) NiCl2(aq)

• Check the activity series. Al more reactive.
• If more reactive metal uncombined, then you will
  get reaction.
• Reaction.
• Determine what ion the uncombined metal will
  form. Al Al3
• Determine formula of new compound. AlCl3
• Complete and balance check solubility of salt.

2 Al(s) 3 NiCl2(aq) 3 Ni(s) 2 AlCl3(aq)
Tro's Introductory Chemistry, Chapter 16
47
48
PracticePredict the Products and Balance the
Equation.

• Mg H3PO4
• Cu H2SO4
• Al Fe2

49
PracticePredict the Products and Balance the
Equation, Continued.

• 3 Mg 2 H3PO4 Mg3(PO4)2 3 H2
• Cu H2SO4 No reaction.
• 2 Al 3 Fe2 2 Al3 3 Fe

50
Electrical Current

• When we talk about the current of a liquid in a
  stream, we are discussing the amount of water
  that passes by in a given period of time.
• When we discuss electric current, we are
  discussing the amount of electric charge that
  passes a point in a given period of time.
• Whether as electrons flowing through a wire or
  ions flowing through a solution.

Tro's Introductory Chemistry, Chapter 16
50
51
Redox Reactions and Current

• Redox reactions involve the transfer of electrons
  from one substance to another.
• Therefore, redox reactions have the potential to
  generate an electric current.
• In order to use that current, we need to separate
  the place where oxidation is occurring from the
  place that reduction is occurring.

Tro's Introductory Chemistry, Chapter 16
51
52
Electric Current Flowing Directly Between Atoms
52
53
Electric Current Flowing Indirectly Between Atoms
Tro's Introductory Chemistry, Chapter 16
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54
Electrochemical Cells

• Electrochemistry is the study of redox reactions
  that produce or require an electric current.
• The conversion between chemical energy and
  electrical energy is carried out in an
  electrochemical cell.
• Spontaneous redox reactions take place in a
  voltaic cell.
• Also known as galvanic cells.
• Batteries are voltaic cells.
• Nonspontaneous redox reactions can be made to
  occur in an electrolytic cell by the addition of
  electrical energy.

55
Electrochemical Cells, Continued

• Oxidation and reduction reactions kept separate.
• Half-cells.
• Electron flow through a wire, along with ion flow
  through a solution, constitutes an electric
  circuit.
• Requires a conductive solid (metal or graphite)
  electrode to allow the transfer of electrons.
• Through external circuit.
• Ion exchange between the two halves of the
  system.
• Electrolyte.

56
Electrodes

• Anode
• Electrode where oxidation occurs.
• Anions attracted to it.
• Connected to positive end of battery in
  electrolytic cell.
• Loses weight in electrolytic cell.
• Cathode
• Electrode where reduction occurs.
• Cations attracted to it.
• Connected to negative end of battery in
  electrolytic cell.
• Gains weight in electrolytic cell.
• Electrode where plating takes place in
  electroplating.

57
Voltaic Cell
58
Current and Voltage

• The number of electrons that flow through the
  system per second is the current.
• Electrode surface area dictates the number of
  electrons that can flow.
• The amount of force pushing the electrons through
  the wire is the voltage.
• The farther the metals are separated on the
  activity series, the larger the voltage will be.

59
Current
The number of electrons that pass a point each
second is called the current of the electricity.
The amount of water that passes a point each
second is called the current of the river.
60
Voltage
Voltage is the force pushing the electrons down
the wire.
Gravity is the force pulling the water down the
river.
61
Dead Battery
As the reaction proceeds, the reactants
get consumed and the voltaic cell dies. The
current decreases until electrons can no
longer flow through the wire.
62
LeClanchés Acidic Dry Cell

• Electrolyte in paste form.
• ZnCl2 NH4Cl.
• Or MgBr2.
• Anode Zn (or Mg).
• Zn(s) Zn2(aq) 2 e-
• Cathode graphite rod.
• MnO2 is reduced.
• 2 MnO2(s) 2 NH4(aq) 2 H2O(l) 2 e- 2
  NH4OH(aq) 2 Mn(O)OH(s)
• Cell voltage 1.5 v.
• Expensive, nonrechargeable, heavy, easily
  corroded.

63
Alkaline Dry Cell

• Same basic cell as acidic dry cell, except
  electrolyte is alkaline KOH paste.
• Anode Zn (or Mg).
• Zn(s) Zn2(aq) 2 e-
• Cathode brass rod.
• MnO2 is reduced.
• 2 MnO2(s) 2 NH4(aq) 2 H2O(l) 2 e- 2
  NH4OH(aq) 2 Mn(O)OH(s)
• Cell voltage 1.54 v.
• Longer shelf life than acidic dry cells and
  rechargeable little corrosion of zinc.

64
Lead Storage Battery

• Six cells in series.
• Electrolyte 6 M H2SO4.
• Anode Pb.
• Pb(s) SO42-(aq) PbSO4(s) 2 e-
• Cathode Pb coated with PbO2.
• PbO2 is reduced.
• PbO2(s) 4 H(aq) SO42-(aq) 2 e- PbSO4(s)
  2 H2O(l)
• Cell voltage 2.09 v.
• Rechargeable, heavy.

65
Fuel Cells

• Like batteries in which reactants are constantly
  being added.
• So it never runs down!
• Anode and cathode both Pt-coated metal.
• Electrolyte is OH solution.
• Anode reaction 2 H2 4 OH ? 4 H2O(l) 4 e-.
• Cathode reaction O2 4 H2O 4 e- ? 4 OH.

66
Nonspontaneous Redox Reaction

• The reverse of a spontaneous reaction is
  nonspontaneous.
• To get it to run, an outside energy source must
  be supplied.
• Nonspontaneous redox reactions can be made to
  work by using a battery to force the electrons to
  flow in the nonspontaneous direction.

67
Electrolysis

• Electrolysis is the process of using electricity
  to break a compound apart.
• Electrolysis is done in an electrolytic cell.
• Electrolytic cells can be used to separate
  elements from their compounds.
• Generate H2 from water for fuel cells.
• Recover metals from their ores.

68
Electrolytic Cell

• The terminal of the battery anode.
• The - terminal of the battery cathode.
• Cations attracted to the cathode anions
  attracted to the anode.
• Cations pick up electrons from the cathode and
  are reduced anions release electrons to the
  anode and are oxidized.
• In electroplating, the work piece is the cathode.
• Cations are reduced at the cathode and plate onto
  the surface.
• The anode is made of the plate metal, the anode
  oxidizes and replaces the metal cations lost from
  the solution.

69
Electrolytic CellElectroplating
70
Corrosion

• Corrosion is the spontaneous oxidation of a metal
  by chemicals in the environment.
• Since many materials we use are active metals,
  corrosion can be a very big problem.

Tro's Introductory Chemistry, Chapter 16
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71
Preventing Corrosion

• One way to reduce or slow corrosion is to coat
  the metal surface to keep it from contacting
  corrosive chemicals in the environment.
• Paint.
• Some metals, like Al, form an oxide that strongly
  attaches to the metal surface, preventing the
  rest from corroding.
• Another method to protect one metal is to attach
  it to a more reactive metal that is cheap.
• Sacrificial electrode.

Tro's Introductory Chemistry, Chapter 16
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