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7/28/2019 ElectroChemistry PPT
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CHEM 160 General Chemistry II
Lecture Presentation
Electrochemistry
Chapter 20
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Electrochemistry
Electrochemistry
deals with interconversion between chemical and
electrical energy
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Electrochemistry
Electrochemistry
deals with the interconversion between chemical and
electrical energy involves redox reactions
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Electrochemistry
Electrochemistry
deals with interconversion between chemical and
electrical energy involves redox reactions
electron transfer reactions
Oh No! Theyre back!
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Redox reactions (quick review)
Oxidation
Reduction
Reducing agent
Oxidizing agent
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Redox reactions (quick review)
Oxidation
loss of electrons
Reduction
Reducing agent
Oxidizing agent
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Redox reactions (quick review)
Oxidation
loss of electrons
Reduction gain of electrons
Reducing agent
Oxidizing agent
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Redox reactions (quick review)
Oxidation
loss of electrons
Reduction gain of electrons
Reducing agent
donates the electrons and is oxidized
Oxidizing agent
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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
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Redox Reactions
Direct redox reaction
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Redox Reactions
Direct redox reaction
Oxidizing and reducing agents are mixed together
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CuSO4(aq)
(Cu2+)
Zn rod
Direct Redox Reaction
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CuSO4(aq)
(Cu2+)
Zn rod
Deposit of
Cu metal
forms
Direct Redox Reaction
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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
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Electrochemical Cells
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Electrochemical Cells
Voltaic Cell
cell in which a spontaneous redox reaction generates
electricity
chemical energy electrical energy
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Electrochemical Cells
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Voltaic Cell
Electrochemical Cells
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Electrochemical Cells
Electrolytic Cell
electrochemical cell in which an electric current
drives a nonspontaneous redox reaction
electrical energy chemical energy
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Cell Potential
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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, Ecell
1 V = 1 J/C
driving force of the redox reaction
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high electrical
potential
low electrical
potential
Cell Potential
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Cell Potential
Ecell = Ecathode - Eanode = Eredn - Eox
Ecell = Ecathode - Eanode = Eredn - Eox
(Ecathode and Eanode are reduction potentials by definition.)
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Cell Potential
Ecell = Ecathode - Eanode = Eredn - Eox
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
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Standard Hydrogen Electrode
E = 0 V (by
definition; arbitrarily
selected)
2H+ + 2e- H2
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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?
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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?
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Standard Electrode Potentials
Standard Reduction Potentials, E
Ecell 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+
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Standard Reduction PotentialsReduction Half-Reaction E(V)F2(g) + 2e
- 2F-(aq) 2.87Au3+(aq) + 3e-Au(s) 1.50Cl2(g) + 2 e
- 2Cl-(aq) 1.36Cr2O7
2-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O 1.33O2(g) + 4H
+ + 4e- 2H2O(l) 1.23Ag+(aq) + e-Ag(s) 0.80Fe3+(aq) + e- Fe2+(aq) 0.77Cu2+(aq) + 2e- Cu(s) 0.34Sn4+(aq) + 2e- Sn2+(aq) 0.152H+(aq) + 2e-H2(g) 0.00Sn2+(aq) + 2e- Sn(s) -0.14Ni2+(aq) + 2e- Ni(s) -0.23Fe2+(aq) + 2e- Fe(s) -0.44Zn2+(aq) + 2e- Zn(s) -0.76Al3+(aq) + 3e-Al(s) -1.66Mg2+(aq) + 2e- Mg(s) -2.37Li+(aq) + e- Li(s) -3.04
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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
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Standard Reduction PotentialsReduction Half-Reaction E(V)F2(g) + 2e
- 2F-(aq) 2.87Au3+(aq) + 3e-Au(s) 1.50Cl2(g) + 2 e
- 2Cl-(aq) 1.36Cr2O7
2-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O 1.33O2(g) + 4H
+ + 4e- 2H2O(l) 1.23Ag+(aq) + e-Ag(s) 0.80Fe3+(aq) + e- Fe2+(aq) 0.77Cu2+(aq) + 2e- Cu(s) 0.34Sn4+(aq) + 2e- Sn2+(aq) 0.152H+(aq) + 2e-H2(g) 0.00Sn2+(aq) + 2e- Sn(s) -0.14Ni2+(aq) + 2e- Ni(s) -0.23Fe2+(aq) + 2e- Fe(s) -0.44Zn2+(aq) + 2e- Zn(s) -0.76Al3+(aq) + 3e-Al(s) -1.66Mg2+(aq) + 2e- Mg(s) -2.37Li+(aq) + e- Li(s) -3.04
Ox.ag
entstrengthincreases
R
ed.agentstre
ngthincrease
s
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Uses of Standard Reduction
Potentials
Determine if oxidizing and reducing agent react
spontaneously
diagonal ruleox. agent
red. agent
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Uses of Standard Reduction
Potentials
Determine if oxidizing and reducing agent react
spontaneously
Cathode
(reduction)
more +
Anode
(oxidation)
more -
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Standard Reduction PotentialsReduction Half-Reaction E(V)F2(g) + 2e
- 2F-(aq) 2.87Au3+(aq) + 3e-Au(s) 1.50Cl2(g) + 2 e
- 2Cl-(aq) 1.36Cr2O7
2-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O 1.33O2(g) + 4H
+ + 4e- 2H2O(l) 1.23Ag+(aq) + e-Ag(s) 0.80Fe3+(aq) + e- Fe2+(aq) 0.77Cu2+(aq) + 2e- Cu(s) 0.34Sn4+(aq) + 2e- Sn2+(aq) 0.152H+(aq) + 2e-H2(g) 0.00Sn2+(aq) + 2e- Sn(s) -0.14Ni2+(aq) + 2e- Ni(s) -0.23Fe2+(aq) + 2e- Fe(s) -0.44Zn2+(aq) + 2e- Zn(s) -0.76Al3+(aq) + 3e-Al(s) -1.66Mg2+(aq) + 2e- Mg(s) -2.37Li+(aq) + e- Li(s) -3.04
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Uses of Standard Reduction
Potentials
Calculate Ecell Ecell = Ecathode - Eanode
Greater Ecell
, greater the driving force
Ecell > 0 : spontaneous redox reactions
Ecell < 0 : nonspontaeous redox reactions
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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 Ecell
for the spontaneous cell reaction
at 25C.
St d d R d ti P t ti l
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Standard Reduction PotentialsReduction Half-Reaction E(V)F2(g) + 2e
- 2F-(aq) 2.87Au3+(aq) + 3e-Au(s) 1.50Cl2(g) + 2 e
- 2Cl-(aq) 1.36Cr2O7
2-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O 1.33O2(g) + 4H
+ + 4e- 2H2O(l) 1.23Ag+(aq) + e-Ag(s) 0.80Fe3+(aq) + e- Fe2+(aq) 0.77Cu2+(aq) + 2e- Cu(s) 0.34Sn4+(aq) + 2e- Sn2+(aq) 0.152H+(aq) + 2e-H2(g) 0.00Sn2+(aq) + 2e- Sn(s) -0.14Ni2+(aq) + 2e- Ni(s) -0.23Fe2+(aq) + 2e- Fe(s) -0.44Zn2+(aq) + 2e- Zn(s) -0.76Al3+(aq) + 3e-Al(s) -1.66Mg2+(aq) + 2e- Mg(s) -2.37Li+(aq) + e- Li(s) -3.04
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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 Ecell
for the spontaneous cell reaction
at 25C.
St d d R d ti P t ti l
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Standard Reduction PotentialsReduction Half-Reaction E(V)F2(g) + 2e
- 2F-(aq) 2.87Au3+(aq) + 3e-Au(s) 1.50Cl2(g) + 2 e
- 2Cl-(aq) 1.36Cr2O7
2-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O 1.33O2(g) + 4H
+ + 4e- 2H2O(l) 1.23Ag+(aq) + e-Ag(s) 0.80Fe3+(aq) + e- Fe2+(aq) 0.77Cu2+(aq) + 2e- Cu(s) 0.34Sn4+(aq) + 2e- Sn2+(aq) 0.152H+(aq) + 2e-H2(g) 0.00Sn2+(aq) + 2e- Sn(s) -0.14Ni2+(aq) + 2e- Ni(s) -0.23Fe2+(aq) + 2e- Fe(s) -0.44Zn2+(aq) + 2e- Zn(s) -0.76Al3+(aq) + 3e-Al(s) -1.66Mg2+(aq) + 2e- Mg(s) -2.37Li+(aq) + e- Li(s) -3.04
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Cell Potential
Is there a relationship between Ecell and DG for a
redox reaction?
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Cell Potential
Relationship between Ecell and DG: DG = -nFEcell
F = Faraday constant = 96500 C/mol e-s, n = # e-s
transferred redox rxn.
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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
DG < 0, Ecell > 0 = spontaneous
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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
DG < 0, Ecell > 0 = spontaneous
Under standard state conditions:
DG = -nFEcell
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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
DG < 0, Ecell > 0 = spontaneous
Under standard state conditions:
DG = -nFEcell
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Equilibrium Constants from Ecell
Relationship between Ecell andD
G: DG = -nFEcell
F = Faraday constant = 96500 C/mol e-s, n = # e-s transferred redoxrxn
1 J = CV
DG < 0, Ecell > 0 = spontaneous
Under standard state conditions:
DG = -nFEcell
and
DG = -RTlnK
so
-nFEcell = -RTlnK
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DH DSCalorimetric Data
DGElectrochemical
DataComposition
DataEcell
Equilibrium
constants
K
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Example 5
Calculate Ecell, DG, and K for the voltaic cell
that uses the reaction between Ag and Cl2 under
standard state conditions at 25C.
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The Nernst Equation
DG depends on concentrations DG = DG + RTlnQ
and
DG = -nFEcell and DG = -nFEcellthus
-nFEcell = -nFEcell + RTlnQ
or
Ecell = Ecell - (RT/nF)lnQ (Nernst eqn.)
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The Nernst Equation
Ecell = Ecell - (RT/nF)lnQ (Nernst eqn.)
At 298 K (25C), RT/F = 0.0257 V
so Ecell = Ecell - (0.0257/n)lnQ
or
Ecell = Ecell - (0.0592/n)logQ
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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 Potentials
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Standard Reduction PotentialsReduction Half-Reaction E(V)F2(g) + 2e
- 2F-(aq) 2.87Au3+(aq) + 3e-Au(s) 1.50Cl2(g) + 2 e
- 2Cl-(aq) 1.36Cr2O7
2-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O 1.33O2(g) + 4H
+ + 4e- 2H2O(l) 1.23Ag+(aq) + e-Ag(s) 0.80Fe3+(aq) + e- Fe2+(aq) 0.77Cu2+(aq) + 2e- Cu(s) 0.34Sn4+(aq) + 2e- Sn2+(aq) 0.152H+(aq) + 2e-H2(g) 0.00Sn2+(aq) + 2e- Sn(s) -0.14Ni2+(aq) + 2e- Ni(s) -0.23Fe2+(aq) + 2e- Fe(s) -0.44Zn2+(aq) + 2e- Zn(s) -0.76Al3+(aq) + 3e-Al(s) -1.66Mg2+(aq) + 2e- Mg(s) -2.37Li+(aq) + e- Li(s) -3.04
Ox.agentstrength
increases
R
ed.agentstre
ngthincrease
s
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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
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How Does a Battery Work
cathode (+)
anode (-)
Electrolyte
Paste
Seal/cap
Assume a generalized battery
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Battery
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 oncompletes the circuit and allows
electron flow to occur
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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 isreversed (nonspontaneous)
Recharge: O2 + R1 R2 + O1
Performed using electrical energy from an externalpower source
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Batteries
Read the textbook to fill in the details on
specific batteries.
Alkaline battery
Lead storage battery
Nicad battery
Fuel cell
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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 Iron
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Corrosion of Iron
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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 isspontaneous!
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Corrosion of Iron
Rust formation:4Fe2+(aq) + O2(g) + 4H
+(aq) 4Fe3+(aq) + 2H2O(l)
2Fe3+(aq) + 4H2O(l) Fe2O3H2O(s) + 6H+(aq)
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Prevention of Corrosion
Cover the Fe surface with a protective coating Paint
Passivation
surface atoms made inactive via oxidation2Fe(s) + 2Na2CrO4(aq) + 2H2O(l) -->
Fe2O3(s) + Cr2O3(s) + 4NaOH(aq)
Other metal
Tin
Zn
Galvanized iron
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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 Potentials
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Standard Electrode Potentials
Half-reaction EF
2(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 havemore negative
Es
Cathodic Protection
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Cathodic Protection
galvanized steel (Fe)
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Cathodic Protection
(cathode)
(electrolyte)
(anode)
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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
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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 theelectrons through the external circuit
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Electrolysis of Molten NaCl
Q tit ti A t f El t h i l C ll
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Quantitative Aspects of Electrochemical Cells
For any half-reaction, the amount of a substanceoxidized or reduced at an electrode is proportional to
the number of electrons passed through the cell
Faradays 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 haspassed
1 C = 1 A x 1 s
1 F = 1 mole e- = 96500 C
Steps for Quantitative Electrolysis
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Steps for Quantitative Electrolysis
Calculations
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
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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?
E l 9
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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.