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CHEM 160 General Chemistry II Lecture Presentation Electrochemistry. December 1, 2004 Chapter 20. Electrochemistry. Electrochemistry deals with interconversion between chemical and electrical energy. Electrochemistry. Electrochemistry

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chem 160 general chemistry ii lecture presentation electrochemistry

CHEM 160 General Chemistry IILecture PresentationElectrochemistry

December 1, 2004

Chapter 20

electrochemistry
Electrochemistry
  • Electrochemistry
    • deals with interconversion between chemical and electrical energy
electrochemistry3
Electrochemistry
  • Electrochemistry
    • deals with the interconversion between chemical and electrical energy
    • involves redox reactions
electrochemistry4
Electrochemistry
  • Electrochemistry
    • deals with interconversion between chemical and electrical energy
    • involves redox reactions
      • electron transfer reactions
      • Oh No! They’re back!
redox reactions quick review
Redox reactions (quick review)
  • Oxidation
  • Reduction
  • Reducing agent
  • Oxidizing agent
redox reactions quick review6
Redox reactions (quick review)
  • Oxidation
    • loss of electrons
  • Reduction
  • Reducing agent
  • Oxidizing agent
redox reactions quick review7
Redox reactions (quick review)
  • Oxidation
    • loss of electrons
  • Reduction
    • gain of electrons
  • Reducing agent
  • Oxidizing agent
redox reactions quick review8
Redox reactions (quick review)
  • Oxidation
    • loss of electrons
  • Reduction
    • gain of electrons
  • Reducing agent
    • donates the electrons and is oxidized
  • Oxidizing agent
redox reactions quick review9
Redox reactions (quick review)
  • Oxidation
    • loss of electrons
  • Reduction
    • gain of electrons
  • Reducing agent
    • donates the electrons and is oxidized
  • Oxidizing agent
    • accepts electrons and is reduced
redox reactions
Redox Reactions
  • Direct redox reaction
redox reactions11
Redox Reactions
  • Direct redox reaction
    • Oxidizing and reducing agents are mixed together
slide12

Direct Redox Reaction

Zn rod

CuSO4(aq) (Cu2+)

slide13

Direct Redox Reaction

Zn rod

CuSO4(aq) (Cu2+)

Deposit of Cu metal forms

redox reactions14
Redox Reactions
  • Direct redox reaction
    • Oxidizing and reducing agents are mixed together
  • Indirect redox reaction
    • Oxidizing and reducing agents are separated but connected electrically
      • Example
        • Zn and Cu2+ can be reacted indirectly
    • Basis for electrochemistry
        • Electrochemical cell
electrochemical cells
Electrochemical Cells
  • Voltaic Cell
    • cell in which a spontaneous redox reaction generates electricity
    • chemical energy  electrical energy
electrochemical cells19
Electrochemical Cells
  • Electrolytic Cell
    • electrochemical cell in which an electric current drives a nonspontaneous redox reaction
    • electrical energy  chemical energy
cell potential21
Cell Potential
  • Cell Potential (electromotive force), Ecell (V)
    • electrical potential difference between the two electrodes or half-cells
      • Depends on specific half-reactions, concentrations, and temperature
      • Under standard state conditions ([solutes] = 1 M, Psolutes = 1 atm), emf = standard cell potential, Ecell
      • 1 V = 1 J/C
    • driving force of the redox reaction
slide22

Cell Potential

low electrical potential

high electrical potential

cell potential23
Cell Potential

Ecell = Ecathode - Eanode = Eredn - Eox

E°cell = E°cathode - E°anode = E°redn - E°ox

(Ecathode and Eanode are reduction potentials by definition.)

cell potential24
Cell Potential
  • E°cell = E°cathode - E°anode = E°redn - E°ox
    • Ecell can be measured
      • Absolute Ecathode and Eanode values cannot
  • Reference electrode
    • has arbitrarily assigned E
    • used to measure relative Ecathode and Eanode for half-cell reactions
  • Standard hydrogen electrode (S.H.E.)
    • conventional reference electrode
standard hydrogen electrode
Standard Hydrogen Electrode
  • E = 0 V (by definition; arbitrarily selected)
  • 2H+ + 2e-  H2
example 1
Example 1

A voltaic cell is made by connecting a standard Cu/Cu2+ electrode to a S.H.E. The cell potential is 0.34 V. The Cu electrode is the cathode. What is the standard reduction potential of the Cu/Cu2+ electrode?

example 2
Example 2

A voltaic cell is made by connecting a standard Zn/Zn2+ electrode to a S.H.E. The cell potential is 0.76 V. The Zn electrode is the anode of the cell. What is the standard reduction potential of the Zn/Zn2+ electrode?

standard electrode potentials
Standard Electrode Potentials
  • Standard Reduction Potentials, E°
    • E°cell measured relative to S.H.E. (0 V)
      • electrode of interest = cathode
    • If E° < 0 V:
      • Oxidizing agent is harder to reduce than H+
    • If E° > 0 V:
      • Oxidizing agent is easier to reduce than H+
standard reduction potentials

Reduction Half-Reaction

E(V)

F2(g) + 2e- 2F-(aq)

2.87

Au3+(aq) + 3e- Au(s)

1.50

Cl2(g) + 2 e- 2Cl-(aq)

1.36

Cr2O72-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O

1.33

O2(g) + 4H+ + 4e- 2H2O(l)

1.23

Ag+(aq) + e- Ag(s)

0.80

Fe3+(aq) + e- Fe2+(aq)

0.77

Cu2+(aq) + 2e- Cu(s)

0.34

Sn4+(aq) + 2e- Sn2+(aq)

0.15

2H+(aq) + 2e- H2(g)

0.00

Sn2+(aq) + 2e- Sn(s)

-0.14

Ni2+(aq) + 2e- Ni(s)

-0.23

Fe2+(aq) + 2e- Fe(s)

-0.44

Zn2+(aq) + 2e- Zn(s)

-0.76

Al3+(aq) + 3e- Al(s)

-1.66

Mg2+(aq) + 2e- Mg(s)

-2.37

Li+(aq) + e- Li(s)

-3.04

Standard Reduction Potentials
uses of standard reduction potentials
Uses of Standard Reduction Potentials
  • Compare strengths of reducing/oxidizing agents.
    • the more - E°, stronger the red. agent
    • the more + E°, stronger the ox. agent
standard reduction potentials33

Reduction Half-Reaction

E(V)

F2(g) + 2e- 2F-(aq)

2.87

Au3+(aq) + 3e- Au(s)

1.50

Cl2(g) + 2 e- 2Cl-(aq)

1.36

Cr2O72-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O

1.33

O2(g) + 4H+ + 4e- 2H2O(l)

1.23

Ag+(aq) + e- Ag(s)

0.80

Fe3+(aq) + e- Fe2+(aq)

0.77

Cu2+(aq) + 2e- Cu(s)

0.34

Sn4+(aq) + 2e- Sn2+(aq)

0.15

2H+(aq) + 2e- H2(g)

0.00

Sn2+(aq) + 2e- Sn(s)

-0.14

Ni2+(aq) + 2e- Ni(s)

-0.23

Fe2+(aq) + 2e- Fe(s)

-0.44

Zn2+(aq) + 2e- Zn(s)

-0.76

Al3+(aq) + 3e- Al(s)

-1.66

Mg2+(aq) + 2e- Mg(s)

-2.37

Li+(aq) + e- Li(s)

-3.04

Standard Reduction Potentials

Ox. agent strength increases

Red. agent strength increases

uses of standard reduction potentials34
Uses of Standard Reduction Potentials
  • Determine if oxidizing and reducing agent react spontaneously
    • diagonal rule

ox. agent

spontaneous

red. agent

uses of standard reduction potentials35
Uses of Standard Reduction Potentials
  • Determine if oxidizing and reducing agent react spontaneously

more +

Spontaneous rxn if E°cathode > E°anode

Cathode (reduction)

E°redn (cathode)

E°redn (V)

Anode (oxidation)

E°redn (anode)

more -

standard reduction potentials36

Reduction Half-Reaction

E(V)

F2(g) + 2e- 2F-(aq)

2.87

Au3+(aq) + 3e- Au(s)

1.50

Cl2(g) + 2 e- 2Cl-(aq)

1.36

Cr2O72-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O

1.33

O2(g) + 4H+ + 4e- 2H2O(l)

1.23

Ag+(aq) + e- Ag(s)

0.80

Fe3+(aq) + e- Fe2+(aq)

0.77

Cu2+(aq) + 2e- Cu(s)

0.34

Sn4+(aq) + 2e- Sn2+(aq)

0.15

2H+(aq) + 2e- H2(g)

0.00

Sn2+(aq) + 2e- Sn(s)

-0.14

Ni2+(aq) + 2e- Ni(s)

-0.23

Fe2+(aq) + 2e- Fe(s)

-0.44

Zn2+(aq) + 2e- Zn(s)

-0.76

Al3+(aq) + 3e- Al(s)

-1.66

Mg2+(aq) + 2e- Mg(s)

-2.37

Li+(aq) + e- Li(s)

-3.04

Standard Reduction Potentials
uses of standard reduction potentials37
Uses of Standard Reduction Potentials
  • Calculate E°cell
    • E°cell = E°cathode - E°anode
      • Greater E°cell, greater the driving force
    • E°cell > 0 : spontaneous redox reactions
    • E°cell < 0 : nonspontaeous redox reactions
example 3
Example 3

A voltaic cell consists of a Ag electrode in 1.0 M AgNO3 and a Cu electrode in 1 M Cu(NO3)2. Calculate E°cell for the spontaneous cell reaction at 25°C.

standard reduction potentials39

Reduction Half-Reaction

E(V)

F2(g) + 2e- 2F-(aq)

2.87

Au3+(aq) + 3e- Au(s)

1.50

Cl2(g) + 2 e- 2Cl-(aq)

1.36

Cr2O72-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O

1.33

O2(g) + 4H+ + 4e- 2H2O(l)

1.23

Ag+(aq) + e- Ag(s)

0.80

Fe3+(aq) + e- Fe2+(aq)

0.77

Cu2+(aq) + 2e- Cu(s)

0.34

Sn4+(aq) + 2e- Sn2+(aq)

0.15

2H+(aq) + 2e- H2(g)

0.00

Sn2+(aq) + 2e- Sn(s)

-0.14

Ni2+(aq) + 2e- Ni(s)

-0.23

Fe2+(aq) + 2e- Fe(s)

-0.44

Zn2+(aq) + 2e- Zn(s)

-0.76

Al3+(aq) + 3e- Al(s)

-1.66

Mg2+(aq) + 2e- Mg(s)

-2.37

Li+(aq) + e- Li(s)

-3.04

Standard Reduction Potentials
example 4
Example 4

A voltaic cell consists of a Ni electrode in 1.0 M Ni(NO3)2 and an Fe electrode in 1 M Fe(NO3)2. Calculate E°cell for the spontaneous cell reaction at 25°C.

standard reduction potentials41

Reduction Half-Reaction

E(V)

F2(g) + 2e- 2F-(aq)

2.87

Au3+(aq) + 3e- Au(s)

1.50

Cl2(g) + 2 e- 2Cl-(aq)

1.36

Cr2O72-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O

1.33

O2(g) + 4H+ + 4e- 2H2O(l)

1.23

Ag+(aq) + e- Ag(s)

0.80

Fe3+(aq) + e- Fe2+(aq)

0.77

Cu2+(aq) + 2e- Cu(s)

0.34

Sn4+(aq) + 2e- Sn2+(aq)

0.15

2H+(aq) + 2e- H2(g)

0.00

Sn2+(aq) + 2e- Sn(s)

-0.14

Ni2+(aq) + 2e- Ni(s)

-0.23

Fe2+(aq) + 2e- Fe(s)

-0.44

Zn2+(aq) + 2e- Zn(s)

-0.76

Al3+(aq) + 3e- Al(s)

-1.66

Mg2+(aq) + 2e- Mg(s)

-2.37

Li+(aq) + e- Li(s)

-3.04

Standard Reduction Potentials
cell potential42
Cell Potential
  • Is there a relationship between Ecell and DG for a redox reaction?
cell potential43
Cell Potential
  • Relationship between Ecell and DG:
    • DG = -nFEcell
      • F = Faraday constant = 96500 C/mol e-’s, n = # e-’s transferred redox rxn.
cell potential44
Cell Potential
  • Relationship between Ecell and DG:
    • DG = -nFEcell
      • F = Faraday constant = 96500 C/mol e-’s, n = # e-’s transferred redox rxn.
      • 1 J = CV
      • G < 0, Ecell > 0 = spontaneous
equilibrium constants from e cell
Equilibrium Constants from Ecell
  • Relationship between Ecell and DG:
    • DG = -nFEcell
      • F = Faraday constant = 96500 C/mol e-’s, n = # e-’s transferred redox rxn
      • 1 J = CV
      • G < 0, Ecell > 0 = spontaneous
  • Under standard state conditions:
    • DG° = -nFE°cell
equilibrium constants from e cell46
Equilibrium Constants from Ecell
  • Relationship between Ecell and DG:
    • DG = -nFEcell
      • F = Faraday constant = 96500 C/mol e-’s, n = # e-’s transferred redox rxn
      • 1 J = CV
      • G < 0, Ecell > 0 = spontaneous
  • Under standard state conditions:
    • DG° = -nFE°cell
equilibrium constants from e cell47
Equilibrium Constants from Ecell
  • Relationship between Ecell and DG:
    • DG = -nFEcell
      • F = Faraday constant = 96500 C/mol e-’s, n = # e-’s transferred redox rxn
      • 1 J = CV
      • G < 0, Ecell > 0 = spontaneous
  • Under standard state conditions:
    • DG° = -nFE°cell

and

    • DG° = -RTlnK

so

  • -nFE°cell = -RTlnK
slide48

Calorimetric Data

DH°

DS°

Electrochemical Data

Composition Data

DG°

E°cell

Equilibrium constants

K

example 5
Example 5

Calculate E°cell, DG°, and K for the voltaic cell that uses the reaction between Ag and Cl2 under standard state conditions at 25°C.

the nernst equation
The Nernst Equation
  • DG depends on concentrations
    • DG = DG° + RTlnQ

and

  • DG = -nFEcell and DG° = -nFE°cell

thus

  • -nFEcell = -nFE°cell + RTlnQ

or

  • Ecell = E°cell - (RT/nF)lnQ (Nernst eqn.)
the nernst equation51
The Nernst Equation
  • Ecell = E°cell - (RT/nF)lnQ (Nernst eqn.)
    • At 298 K (25°C), RT/F = 0.0257 V

so

  • Ecell = E°cell - (0.0257/n)lnQ

or

  • Ecell = E°cell - (0.0592/n)logQ
example 7
Example 7
  • Calculate the voltage produced by the galvanic cell which uses the reaction below if [Ag+] = 0.001 M and [Cu2+] = 1.3 M.

2Ag+(aq) + Cu(s)  2Ag(s) + Cu2+(aq)

standard reduction potentials53

Reduction Half-Reaction

E(V)

F2(g) + 2e- 2F-(aq)

2.87

Au3+(aq) + 3e- Au(s)

1.50

Cl2(g) + 2 e- 2Cl-(aq)

1.36

Cr2O72-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O

1.33

O2(g) + 4H+ + 4e- 2H2O(l)

1.23

Ag+(aq) + e- Ag(s)

0.80

Fe3+(aq) + e- Fe2+(aq)

0.77

Cu2+(aq) + 2e- Cu(s)

0.34

Sn4+(aq) + 2e- Sn2+(aq)

0.15

2H+(aq) + 2e- H2(g)

0.00

Sn2+(aq) + 2e- Sn(s)

-0.14

Ni2+(aq) + 2e- Ni(s)

-0.23

Fe2+(aq) + 2e- Fe(s)

-0.44

Zn2+(aq) + 2e- Zn(s)

-0.76

Al3+(aq) + 3e- Al(s)

-1.66

Mg2+(aq) + 2e- Mg(s)

-2.37

Li+(aq) + e- Li(s)

-3.04

Standard Reduction Potentials

Ox. agent strength increases

Red. agent strength increases

commercial voltaic cells
Commercial Voltaic Cells
  • Battery
    • commercial voltaic cell used as portable source of electrical energy
  • types
    • primary cell
      • Nonrechargeable
      • Example: Alkaline battery
    • secondary cell
      • Rechargeable
      • Example: Lead storage battery
how does a battery work

Seal/cap

cathode (+)

Electrolyte Paste

anode (-)

How Does a Battery Work

Assume a generalized battery

battery
Battery

Placing the battery into a flashlight, etc., and turning the power on completes the circuit and allows electron flow to occur

Electrolyte paste: ion migration occurs here

e- flow

cathode (+): Reduction occurs here

anode (-): oxidation occurs here

how does a battery work57
How Does a Battery Work
  • Battery reaction when producing electricity (spontaneous):

Cathode: O1 + e- R1

Anode: R2  O2 + e-

Overall: O1 + R2  R1 + O2

  • Recharging a secondary cell
    • Redox reaction must be reversed, i.e., current is reversed (nonspontaneous)

Recharge: O2 + R1  R2 + O1

    • Performed using electrical energy from an external power source
batteries
Batteries
  • Read the textbook to fill in the details on specific batteries.
    • Alkaline battery
    • Lead storage battery
    • Nicad battery
    • Fuel cell
corrosion
Corrosion
  • Corrosion
    • deterioration of metals by a spontaneous redox reaction
      • Attacked by species in environment
        • Metal becomes a “voltaic” cell
      • Metal is often lost to a solution as an ion
  • Rusting of Iron
corrosion of iron62
Corrosion of Iron

Half-reactions

anode: Fe(s)  Fe2+(aq) + 2e-

cathode: O2(g) + 4H+(aq) + 4e- 2H2O(l)

overall: 2Fe(s) + O2(g) + 4H+(aq)  2Fe2+(aq) + 2H2O(l)

Ecell > 0 (Ecell = 0.8 to 1.2 V), so process is spontaneous!

corrosion of iron63
Corrosion of Iron

Rust formation:

4Fe2+(aq) + O2(g) + 4H+(aq)  4Fe3+(aq) + 2H2O(l)

2Fe3+(aq) + 4H2O(l)  Fe2O3·H2O(s) + 6H+(aq)

prevention of corrosion
Prevention of Corrosion
  • Cover the Fe surface with a protective coating
    • Paint
    • Passivation
      • surface atoms made inactive via oxidation

2Fe(s) + 2Na2CrO4(aq) + 2H2O(l) --> Fe2O3(s) + Cr2O3(s) + 4NaOH(aq)

    • Other metal
      • Tin
      • Zn
        • Galvanized iron
prevention of corrosion65
Prevention of Corrosion
  • Cathodic Protection
    • metal to be protected is brought into contact with a more easily oxidized metal
    • “sacrificial” metal becomes the anode
      • “Corrodes” preferentially over the iron
      • Iron serves only as the cathode
standard electrode potentials66
Standard Electrode Potentials

Half-reaction E°

F2(g) + 2e- -> 2F-(aq) +2.87 V

Ag+(aq) + e- -> Ag(s) +0.80 V

Cu2+(aq) + 2e- -> Cu(s) +0.34 V

2H+(aq) + 2e- -> H2(g) 0 V

Ni2+(aq) + 2e- -> Ni(s) -0.25 V

Fe2+(aq) + 2e- -> Fe(s) -0.44 V

Zn2+(aq) + 2e- -> Zn(s) -0.76 V

Al3+(aq) + 3e- -> Al(s) -1.66 V

Mg2+(aq) + 2e- ->Mg(s) -2.38 V

Metals more easily oxidized than Fe have more negative E°’s

cathodic protection
Cathodic Protection

galvanized steel (Fe)

cathodic protection68
Cathodic Protection

(anode)

(cathode)

(electrolyte)

electrolysis
Electrolysis
  • Electrolysis
    • process in which electrical energy drives a nonspontaneous redox reaction
      • electrical energy is converted into chemical energy
  • Electrolytic cell
    • electrochemical cell in which an electric current drives a nonspontaneous redox reaction
electrolysis70
Electrolysis
  • Same principles apply to both electrolytic and voltaic cells
    • oxidation occurs at the anode
    • reduction occurs at the cathode
    • electrons flow from anode to cathode in the external circuit
      • In an electrolytic cell, an external power source pumps the electrons through the external circuit
quantitative aspects of electrochemical cells
Quantitative Aspects of Electrochemical Cells
  • For any half-reaction, the amount of a substance oxidized or reduced at an electrode is proportional to the number of electrons passed through the cell
    • Faraday’s law of electrolysis
    • Examples
      • Na+ + 1e- Na
      • Al3+ + 3e-  Al
    • Number of electrons passing through cell is measured by determining the quantity of charge (coulombs) that has passed
      • 1 C = 1 A x 1 s
      • 1 F = 1 mole e- = 96500 C
steps for quantitative electrolysis calculations

current (A) and time (s), A x s

Steps for Quantitative Electrolysis Calculations

charge in coulombs (C)

Number of moles of e-

moles of substance oxidized or reduced

mass of substance oxidized or reduced

example 8
Example 8
  • What mass of copper metal can be produced by a 3.00 A current flowing through a copper(II) sulfate (CuSO4) solution for 5.00 hours?
example 9
Example 9
  • An aqueous solution of an iron salt is electrolyzed by passing a current of 2.50 A for 3.50 hours. As a result, 6.1 g of iron metal are formed at the cathode. Calculate the charge on the iron ions in the solution.