Chapter 18 Oxidation–Reduction Reactions and Electrochemistry

Chapter 18

Oxidation–Reduction Reactions and

Electrochemistry

Chapter 18

Table of Contents

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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

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• 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

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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

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• 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

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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

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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

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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

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More Practice!

Oxidation–Reduction Reactions Between Nonmetals

Section 18.3

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• 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

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• 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

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• 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

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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

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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

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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

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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

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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

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• 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

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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

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• 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

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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

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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

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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

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Making an Electrochemical Cell

8H+ + MnO4– + 5e– → Mn2+ + 4H2O

Fe2+ → Fe3+ + e–

Section 18.5

Electrochemistry: An Introduction

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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

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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

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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

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Electrolysis

• Process where electrical energy is used to produce a chemical change. Nonspontaneous

Section 18.6

Batteries

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Lead Storage Battery

• Anode reaction – oxidationPb + H2SO4 PbSO4 + 2H+ + 2e

• Cathode reaction – reductionPbO2 + H2SO4 + 2e + 2H+ PbSO4 + 2H2O

Section 18.6

Batteries

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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

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Electric Potential

• The “pressure” on electrons to flow from anode to cathode in a battery, like water flow.

Section 18.6

Batteries

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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

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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

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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

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• 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

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• Forcing a current through a cell to produce a chemical change that would not otherwise occur.