Chapter 9 – Chem 160 Chemical Calculations: The Mole Concept and Chemical Formulas.
CHEM 160 General Chemistry II Lecture Presentation Electrochemistry December 1, 2004 Chapter 20.
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Transcript of CHEM 160 General Chemistry II Lecture Presentation Electrochemistry December 1, 2004 Chapter 20.
CHEM 160 General Chemistry IICHEM 160 General Chemistry IILecture PresentationLecture Presentation
ElectrochemistryElectrochemistry
December 1, 2004
Chapter 20
ElectrochemistryElectrochemistry
Electrochemistry deals with interconversion between chemical and
electrical energy
ElectrochemistryElectrochemistry
Electrochemistry deals with the interconversion between chemical and
electrical energy involves redox reactions
ElectrochemistryElectrochemistry
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 review)Redox reactions (quick review)
Oxidation loss of electrons
Reduction
Reducing agent
Oxidizing agent
Redox reactions (quick review)Redox reactions (quick review)
Oxidation loss of electrons
Reduction gain of electrons
Reducing agent
Oxidizing agent
Redox reactions (quick review)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 review)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 ReactionsRedox Reactions
Direct redox reaction
Redox ReactionsRedox Reactions
Direct redox reaction Oxidizing and reducing agents are mixed together
CuSO4(aq) (Cu2+)
Zn rod
Direct Redox ReactionDirect Redox Reaction
CuSO4(aq) (Cu2+)
Zn rod
Deposit of Cu metal
forms
Direct Redox ReactionDirect Redox Reaction
Redox ReactionsRedox 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 CellsElectrochemical Cells
Electrochemical CellsElectrochemical Cells
Voltaic Cell cell in which a spontaneous redox reaction generates
electricity chemical energy electrical energy
Electrochemical CellsElectrochemical Cells
Voltaic Cell
Electrochemical CellsElectrochemical Cells
Electrochemical CellsElectrochemical Cells
Electrolytic Cell electrochemical cell in which an electric current
drives a nonspontaneous redox reaction electrical energy chemical energy
Cell PotentialCell Potential
Cell PotentialCell 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, Ecell
• 1 V = 1 J/C
driving force of the redox reaction
high electrical high electrical potentialpotential
low electrical low electrical potentialpotential
Cell PotentialCell Potential
Cell PotentialCell 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 PotentialCell 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 reactionsStandard hydrogen electrode (S.H.E.)
conventional reference electrode
Standard Hydrogen ElectrodeStandard Hydrogen Electrode
E = 0 V (by definition; arbitrarily selected)
2H+ + 2e- H2
Example 1Example 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 2Example 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 PotentialsStandard 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 PotentialsStandard Reduction PotentialsReduction 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
Uses of Standard Reduction Uses of Standard Reduction PotentialsPotentials
Compare strengths of reducing/oxidizing agents. the more - E°, stronger the red. agent the more + E°, stronger the ox. agent
Standard Reduction PotentialsStandard Reduction PotentialsReduction 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
Ox.
age
nt s
tren
gth
incr
ease
sR
ed. agent strength increases
Uses of Standard Reduction Uses of Standard Reduction PotentialsPotentials
Determine if oxidizing and reducing agent react spontaneously diagonal rule
ox. agent
red. agent
spontaneous
spontaneous
Uses of Standard Reduction Uses of Standard Reduction PotentialsPotentials
Determine if oxidizing and reducing agent react spontaneously
Cathode (reduction) E°redn (cathode)
more +
Anode (oxidation)
E° re
dn (
V)
E°redn (anode)
more -
Spontaneous rxn if Spontaneous rxn if EE°°cathodecathode > E > E°°anodeanode
Standard Reduction PotentialsStandard Reduction PotentialsReduction 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
Uses of Standard Reduction Uses of Standard Reduction PotentialsPotentials
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 3Example 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 PotentialsStandard Reduction PotentialsReduction 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
Example 4Example 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 PotentialsStandard Reduction PotentialsReduction 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
Cell PotentialCell Potential
Is there a relationship between Ecell and G for a redox reaction?
Cell PotentialCell Potential
Relationship between Ecell and G:
G = -nFEcell
• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s transferred redox rxn.
Cell PotentialCell Potential
Relationship between Ecell and G:
G = -nFEcell
• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s transferred redox rxn.
• 1 J = CVG < 0, Ecell > 0 = spontaneous
Equilibrium Constants from EEquilibrium Constants from Ecellcell
Relationship between Ecell and G:
G = -nFEcell
• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s transferred redox rxn
• 1 J = CVG < 0, Ecell > 0 = spontaneous
Under standard state conditions: G° = -nFE°cell
Equilibrium Constants from EEquilibrium Constants from Ecellcell
Relationship between Ecell and G:
G = -nFEcell
• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s transferred redox rxn
• 1 J = CVG < 0, Ecell > 0 = spontaneous
Under standard state conditions: G° = -nFE°cell
Equilibrium Constants from EEquilibrium Constants from Ecellcell
Relationship between Ecell and G: G = -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: G° = -nFE°cell
and G° = -RTlnK
so -nFE°cell = -RTlnK
H° S°
Calorimetric Data
G°Electrochemical
DataComposition
Data
E°cell
Equilibrium constants
K
Example 5Example 5
Calculate E°cell, G°, and K for the voltaic cell that uses the reaction between Ag and Cl2 under standard state conditions at 25°C.
The Nernst EquationThe Nernst EquationG depends on concentrations
G = G° + RTlnQ
andG = -nFEcell and G° = -nFE°cell
thus-nFEcell = -nFE°cell + RTlnQ
or Ecell = E°cell - (RT/nF)lnQ (Nernst eqn.)
The Nernst EquationThe Nernst Equation
Ecell = E°cell - (RT/nF)lnQ (Nernst eqn.) At 298 K (25°C), RT/F = 0.0257 V
soEcell = E°cell - (0.0257/n)lnQ
orEcell = E°cell - (0.0592/n)logQ
Example 7Example 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 PotentialsStandard Reduction PotentialsReduction 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
Ox.
age
nt s
tren
gth
incr
ease
sR
ed. agent strength increases
Commercial Voltaic CellsCommercial Voltaic CellsBattery
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 WorkHow Does a Battery Work
cathode (+)
anode (-)
Electrolyte Paste
Seal/cap
Assume a generalized battery
BatteryBattery
cathode (+): Reduction occurs
here
anode (-): oxidation
occurs here
e- flow
Electrolyte paste: ion migration occurs
here
Placing the battery into a flashlight, etc., and turning the power on completes the circuit and allows
electron flow to occur
How Does a Battery WorkHow Does a Battery WorkBattery 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
BatteriesBatteries
Read the textbook to fill in the details on specific batteries. Alkaline battery Lead storage battery Nicad battery Fuel cell
CorrosionCorrosionCorrosion
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 IronCorrosion of Iron
Corrosion of IronCorrosion 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 IronCorrosion of Iron
Rust formation:
4Fe2+(aq) + O2(g) + 4H+(aq) 4Fe3+(aq) + 2H2O(l)
2Fe3+(aq) + 4H2O(l) Fe2O3H2O(s) + 6H+(aq)
Prevention of CorrosionPrevention 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 CorrosionPrevention 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 PotentialsStandard 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 ProtectionCathodic Protection
galvanized steel (Fe)
Cathodic ProtectionCathodic Protection
(cathode)
(electrolyte)
(anode)
ElectrolysisElectrolysis
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
ElectrolysisElectrolysis
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
Electrolysis of Molten NaClElectrolysis of Molten NaCl
Quantitative Aspects of Electrochemical CellsQuantitative 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 Steps for Quantitative Electrolysis CalculationsCalculations
current (A) and time (s), A x s
charge in coulombs
(C)
Number of moles of e-
moles of substance oxidized or reduced
mass of substance oxidized or reduced
Example 8Example 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 9Example 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.