Upload
leonard-fisher
View
239
Download
3
Embed Size (px)
Citation preview
Chapter 18
Oxidation–Reduction Reactions and
Electrochemistry
Chapter 18
Table of Contents
2
18.1Oxidation–Reduction Reactions
18.2 Oxidation States
18.3 Oxidation–Reduction Reactions Between Nonmetals
18.4Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
18.5Electrochemistry: An Introduction
18.6Batteries
18.7Corrosion
18.8Electrolysis
Section 18.1
Oxidation–Reduction Equations
Return to TOC
3
• Oxidation–reduction reaction (redox reaction) – a chemical reaction involving the transfer of electrons. Oxidation – loss of electrons Reduction – gain of electrons
http://www.youtube.com/watch?v=Ftw7a5ccubs
Section 18.1
Oxidation–Reduction Equations
Return to TOC
Copyright © Cengage Learning. All rights reserved 4
Exercise
In the reaction below Sn(II) _____________.
Sn2+ + 2Fe3+ → Sn4+ + 2Fe2+
a) gains electronsb) is reducedc) is oxidizedd) is neither oxidized nor reduced
Oxidation States
Section 18.2
Return to TOC
5
• Allow us to keep track of electrons in oxidation–reduction reactions by assigning charges to the various atoms in a compound.
Oxidation States for the Transition Metals
Oxidation States
Section 18.2
Return to TOC
Copyright © Cengage Learning. All rights reserved 6
1. Oxidation state of an atom in an elemental state = 0
2. Oxidation state of monatomic ion = charge of the ion
3. Oxygen = 2 in covalent compounds (except in peroxides where it = 1)
4. Hydrogen = +1 in covalent compounds
5. Fluorine = 1 in compounds
6. Sum of oxidation states = 0 in compounds
7. Sum of oxidation states = charge of the ion in ions
Rules for Assigning Oxidation States
Oxidation States
Section 18.2
Return to TOC
Copyright © Cengage Learning. All rights reserved 7
Exercise
Find the oxidation states for each of the elements in each of the following compounds:
• K2Cr2O7
• CO32-
• MnO2
• PCl5• SF4
K = +1; Cr = +6; O = –2
C = +4; O = –2
Mn = +4; O = –2
P = +5; Cl = –1
S = +4; F = –1
Oxidation States
Section 18.2
Return to TOC
Copyright © Cengage Learning. All rights reserved 8
What are the Oxidation Numbers for each element in the following?
H2O
N2
KMnO4
CO2
CH4
CHCl3HeCu
Na2Cr2O7
+1 for H, -2 for O
Zero for N, elemental state+1 for K, -2 for O, +7 for Mn
-2 for O, +4 for C
+1 for H, -4 for C
+1 for H, -1 for Cl, +2 for C
Zero for He, elemental state
Zero for Cu, elemental state
+1 for Na, -2 for O, +6 for Cr
1(+1 K)=+14(-2 O)= -8
-7
1(+1 H)=+13(-1 Cl)= -3 -2
2(+1 Na)=+2 7(-2 O)= -14
-12
Oxidation States
Section 18.2
Return to TOC
9
More Practice!
Oxidation–Reduction Reactions Between Nonmetals
Section 18.3
Return to TOC
10
• 2Na(s) + Cl2(g) 2NaCl(s)
• Na oxidized Na is also called the reducing agent (electron
donor).
• Cl2 reduced
Cl2 is also called the oxidizing agent (electron
acceptor).
Oxidation–Reduction Reactions Between Nonmetals
Section 18.3
Return to TOC
11
• CH4(g) + 2O2(g) CO2(g) + 2H2O(g)
• C oxidized CH4 is the reducing agent.
• O2 reduced
O2 is the oxidizing agent.
Oxidation–Reduction Reactions Between Nonmetals
Section 18.3
Return to TOC
Copyright © Cengage Learning. All rights reserved 12
• Transfer of electrons• Transfer may occur to form ions• Oxidation – increase in oxidation state
(loss of electrons); reducing agent• Reduction – decrease in oxidation state
(gain of electrons); oxidizing agent
Redox Characteristics
0 2+ 1- 2+ 1- 0
Zn(s) + CuCl2(aq) ZnCl2(aq) + Cu(s)
•Reduction
Oxidation
Oxidation–Reduction Reactions Between Nonmetals
Section 18.3
Return to TOC
Copyright © Cengage Learning. All rights reserved 13
Concept Check
Which of the following are oxidation–reduction reactions? Identify the oxidizing agent and the reducing agent.
a)Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g)
b)Cr2O72-(aq) + 2OH-(aq) 2CrO4
2-(aq) + H2O(l)
c)2CuCl(aq) CuCl2(aq) + Cu(s)
Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Return to TOC
Copyright © Cengage Learning. All rights reserved 14
Half–Reactions
• The overall reaction is split into two half–reactions, one involving oxidation and one reduction.
• Has electrons as reactants or products
8H+ + MnO4– + 5Fe2+ → Mn2+ + 5Fe3+ + 4H2O
Reduction: 8H+ + MnO4– + 5e– → Mn2+ + 4H2O
Oxidation: 5Fe2+ → 5Fe3+ + 5e–
Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Return to TOC
Copyright © Cengage Learning. All rights reserved 15
1. Identify and write the equations for the oxidation and reduction half–reactions.
2. For each half–reaction:A. Balance all the elements except H and O.
B. Balance O using H2O.
C. Balance H using H+.
D. Balance the charge using electrons.
The Half–Reaction Method for Balancing Equations for Oxidation–Reduction Reactions Occurring in Acidic Solution
Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Return to TOC
Copyright © Cengage Learning. All rights reserved 16
3. If necessary, multiply one or both balanced half–reactions by an integer to equalize the number of electrons transferred in the two half–reactions.
4. Add the half–reactions, and cancel identical species.
5. Check that the elements and charges are balanced.
The Half–Reaction Method for Balancing Equations for Oxidation–Reduction Reactions Occurring in Acidic Solution
Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Return to TOC
Copyright © Cengage Learning. All rights reserved 17
Cr2O72-(aq) + SO3
2-(aq) Cr3+(aq) + SO42-(aq)
• How can we balance this equation?• First Steps:
Separate into half-reactions. Balance elements except H and O.
Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Return to TOC
Copyright © Cengage Learning. All rights reserved 18
• Cr2O72-(aq) 2Cr3+(aq)
• SO32-(aq) SO4
2-(aq)
• Balance O’s with H2O and H’s with H+
Method of Half Reactions
• 14H+(aq) + Cr2O72-(aq) 2Cr3+(aq) + 7H2O(aq)
• H2O(l) + SO32-(aq) SO4
2-(aq) + 2H+(aq)
• How many electrons are involved in each half reaction? Balance the charges.
Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Return to TOC
Copyright © Cengage Learning. All rights reserved 19
Method of Half Reactions (continued)
6 e- + 14H+(aq) + Cr2O72-(aq) 2Cr3+(aq) + 7H2O(aq)
H2O(l) + SO32-(aq) SO4
2-(aq) + 2H+(aq) + 2e-
Multiply whole reactions by a whole number to make the number of electrons gained equal the number of electrons lost.
6 e- + 14H+(aq) + Cr2O72-(aq) 2Cr3+(aq) + 7H2O(aq)
3(H2O(l) + SO32-(aq) SO4
2-(aq) + 2H+(aq) + 2e-)
Combine half reactions cancelling out those reactants and products that are the same on both sides, especially the electrons.
Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Return to TOC
Copyright © Cengage Learning. All rights reserved 20
• Final Balanced Equation: Cr2O7
2- + 3SO32- + 8H+ 2Cr3+ + 3SO4
2- + 4H2O
Method of Half Reactions (continued)
6e- + 14H+(aq) + Cr2O72-(aq) 2Cr3+(aq) + 7H2O(aq)
3H2O(l) + 3SO32-(aq) 3SO4
2-(aq) + 6H+(aq) + 6e-
48
Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Return to TOC
Copyright © Cengage Learning. All rights reserved 21
Exercise
When the reaction Ce2+ + Co2+ → Ce3+ + Co is balanced, the coefficient in front of Ce2+ is
a) 0b) 1c) 2d) 3
Ce2+ → Ce3+ +1e-
2e- + Co2+ → Co
2Ce2+ + Co2+ → 2Ce3+ + Co
Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Return to TOC
22
Exercise
Balance the following oxidation–reduction reaction that occurs in acidic solution.
Br–(aq) + MnO4–(aq) Br2(l)+ Mn2+(aq)
10Br–(aq) + 16H+(aq) + 2MnO4–(aq) 5Br2(l)+ 2Mn2+(aq) + 8H2O(l)
Section 18.5
Electrochemistry: An Introduction
Return to TOC
23
Electrochemistry
• The study of the interchange of chemical and electrical energy.
• Two types of processes: Production of an electric current from a
chemical reaction. The use of electric current to produce a
chemical change.
Section 18.5
Electrochemistry: An Introduction
Return to TOC
24
Making an Electrochemical Cell
8H+ + MnO4– + 5e– → Mn2+ + 4H2O
Fe2+ → Fe3+ + e–
Section 18.5
Electrochemistry: An Introduction
Return to TOC
Copyright © Cengage Learning. All rights reserved 25
Making an Electrochemical Cell
• If electrons flow through the wire charge builds up.
• Solutions must be connected to permit ions to flow to balance the charge.
Section 18.5
Electrochemistry: An Introduction
Return to TOC
Copyright © Cengage Learning. All rights reserved 26
Making an Electrochemical Cell
• A salt bridge or porous disk connects the half cells and allows ions to flow, completing the circuit.
Section 18.5
Electrochemistry: An Introduction
Return to TOC
Copyright © Cengage Learning. All rights reserved 27
Electrochemical Battery (Galvanic Cell)
• Device powered by an oxidation–reduction reaction where chemical energy is converted to electrical energy.
• Anode – electrode where oxidation occurs • Cathode – electrode where reduction occurs
Section 18.5
Electrochemistry: An Introduction
Return to TOC
Copyright © Cengage Learning. All rights reserved 28
Electrolysis
• Process where electrical energy is used to produce a chemical change. Nonspontaneous
Section 18.6
Batteries
Return to TOC
29
Lead Storage Battery
• Anode reaction – oxidationPb + H2SO4 PbSO4 + 2H+ + 2e
• Cathode reaction – reductionPbO2 + H2SO4 + 2e + 2H+ PbSO4 + 2H2O
Section 18.6
Batteries
Return to TOC
30
Lead Storage Battery – Overall Reaction
Pb(s) + PbO2(s) + 2H2SO4(aq) 2PbSO4(s) + 2H2O(l)
Hydrometer to measure H2SO4
concentration.
As the battery discharges the sulfate of the acid precipitates with the lead taking it out of solution and
reducing the acid concentration. As the battery is recharged the current goes into dissolving the lead
sulfate restoring the acid concentration.
Section 18.6
Batteries
Return to TOC
31
Electric Potential
• The “pressure” on electrons to flow from anode to cathode in a battery, like water flow.
Section 18.6
Batteries
Return to TOC
32
Dry Cell Batteries
• Do not contain a liquid electrolyte.
• Acid version
• Anode reaction – oxidation
Zn Zn2+ + 2e
• Cathode reaction – reduction
2NH4+ + 2MnO2 + 2e Mn2O3 + 2NH3 + 2H2O
Section 18.6
Batteries
Return to TOC
33
Dry Cell Batteries
• Do not contain a liquid electrolyte. Alkaline version– Anode reaction – oxidation
Zn + 2OH ZnO + H2O + 2e
– Cathode reaction – reduction
2MnO2 + H2O + 2e Mn2O3 + 2OH
Section 18.6
Batteries
Return to TOC
Copyright © Cengage Learning. All rights reserved 34
Dry Cell Batteries
• Do not contain a liquid electrolyte. Other Types
• Silver cell – Zn anode, Ag2O cathode
• Mercury cell – Zn anode, HgO cathode
• Nickel-cadmium – rechargeable
Section 18.7
Corrosion
Return to TOC
Copyright © Cengage Learning. All rights reserved 35
• The oxidation of metals to form mainly oxides and sulfides. Some metals, such as aluminum, protect themselves
with their oxide coating. Corrosion of iron can be
prevented by coatings, by alloying and cathodic protection.
Cathodic protection of an underground pipe.
Section 18.8
Electrolysis
Return to TOC
Copyright © Cengage Learning. All rights reserved 36
• Forcing a current through a cell to produce a chemical change that would not otherwise occur.