CHEM 160 General Chemistry II Lecture Presentation Electrochemistry

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CHEM 160 General Chemistry IILecture
PresentationElectrochemistry

• December 1, 2004
• 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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Direct Redox Reaction
Zn rod
CuSO4(aq) (Cu2)
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Direct Redox Reaction
Zn rod
CuSO4(aq) (Cu2)
Deposit of Cu metal forms
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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
Voltaic Cell
Zn anode
Cu cathode
Salt bridge
Cu2
Zn ? Zn2 2e-
Cu2 2e-? Cu
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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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Electrochemical Cells
Voltaic Cell
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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,
  E?cell
• 1 V 1 J/C
• driving force of the redox reaction

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Cell Potential
voltmeter
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Cell Potential
low electrical potential
high electrical 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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H2 (1 atm)
Pt
1 M H
1 M Cu2
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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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e-
e-
H2 (1 atm)
Pt
Zn
1 M H
1 M Zn2
2H 2e- ? H2
Zn ? Zn2 2e-
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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 lt 0 V
• Oxidizing agent is harder to reduce than H
• If E gt 0 V
• Oxidizing agent is easier to reduce than H

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Standard Reduction Potentials
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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 Potentials
Ox. agent strength increases
Red. agent strength increases
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Uses of Standard Reduction Potentials

• Determine if oxidizing and reducing agent react
  spontaneously
• diagonal rule

ox. agent
spontaneous
red. agent
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Uses of Standard Reduction Potentials

• Determine if oxidizing and reducing agent react
  spontaneously

more
Spontaneous rxn if Ecathode gt Eanode
Cathode (reduction)
Eredn (cathode)
Eredn (V)
Anode (oxidation)
Eredn (anode)
more -
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Standard Reduction Potentials
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Uses of Standard Reduction Potentials

• Calculate Ecell
• Ecell Ecathode - Eanode
• Greater Ecell, greater the driving force
• Ecell gt 0 spontaneous redox reactions
• Ecell lt 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.

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Standard Reduction Potentials
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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.

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Standard Reduction Potentials
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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
• ?G lt 0, Ecell gt 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
• ?G lt 0, Ecell gt 0 spontaneous
• Under standard state conditions
• DG -nFEcell

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Example 5

• Calculate DG at room temperature for the
  reaction between Sn4(aq) and Fe(s).

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Standard Reduction Potentials
Ox. agent strength increases
Red. agent strength increases
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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
• ?G lt 0, Ecell gt 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
• ?G lt 0, Ecell gt 0 spontaneous
• Under standard state conditions
• DG -nFEcell
• and
• DG -RTlnK
• so
• -nFEcell -RTlnK

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Calorimetric Data
DH
DS
Electrochemical Data
Composition Data
DG
Ecell
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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Standard Reduction Potentials
Ox. agent strength increases
Red. agent strength increases
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The Nernst Equation

• DG depends on concentrations
• DG DG RTlnQ
• and
• DG -nFEcell and DG -nFEcell
• thus
• -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 voltaic
  cell using the reaction between Zn(s) and
  Cu2(aq) if Zn2 0.001 M and Cu2 1.3 M.
• Zn(s) Cu2(aq) ? Zn2(aq) Cu(s)

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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)

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Standard Reduction Potentials
Ox. agent strength increases
Red. agent strength increases
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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
Assume a generalized battery
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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
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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 is
  reversed (nonspontaneous)
• Recharge O2 R1 ? R2 O1
• Performed using electrical energy from an
  external power 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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Alkaline Dry Cell
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Alkaline Dry Cell
Plated steel ()
Cathode Mixture of MnO2 and C (graphite)
Brass rod
Anode Mixture of Zn and KOH(aq)
Paper or fabric Separator
Insulators
Plated steel (-)
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Alkaline Dry Cell

• Half-reactions

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Alkaline Dry Cell

• Half-reactions
• anode Zn(s) 2OH-(aq) --gt ZnO(s) H2O(l) 2e-

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Alkaline Dry Cell

• Half-reactions
• anode Zn(s) 2OH-(aq) --gt ZnO(s) H2O(l)
  2e-
• cathode 2MnO2(s) H2O(l) 2e- --gt
  Mn2O3(s) 2OH-(aq)

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Alkaline Dry Cell

• Half-reactions
• anode Zn(s) 2OH-(aq) --gt ZnO(s) H2O(l)
  2e-
• cathode 2MnO2(s) H2O(l) 2e- --gt
  Mn2O3(s) 2OH-(aq)
• overall Zn(s) 2MnO2(s) --gt Mn2O3(s) ZnO(s)
• Ecell 1.54 V

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Lead Storage Battery
(anode)
(cathode)
6 x 2V 12 V
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Lead Storage Battery

• Half-reactions

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Lead Storage Battery

• Half-reactions
• anode Pb(s) SO42-(aq) --gt PbSO4(s) 2e-

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Lead Storage Battery

• Half-reactions
• anode Pb(s) SO42-(aq) --gt PbSO4(s) 2e-
• cathode PbO2(s) 4H(aq) SO42-(aq) 2e- --gt
  PbSO4(s) 2H2O(l)

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Lead Storage Battery

• Half-reactions
• anode Pb(s) SO42-(aq) --gt PbSO4(s) 2e-
• cathode PbO2(s) 4H(aq) SO42-(aq) 2e- --gt
  PbSO4(s) 2H2O(l)
• overall Pb(s) PbO2(s) 2H2SO4(aq) --gt
  2PbSO4(s) 2H2O(l)

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Lead Storage Battery

• Half-reactions
• anode Pb(s) SO42-(aq) --gt PbSO4(s) 2e-
• cathode PbO2(s) 4H(aq) SO42-(aq) 2e- --gt
  PbSO4(s) 2H2O(l)
• overall Pb(s) PbO2(s) 2H2SO4(aq) --gt
  2PbSO4(s) 2H2O(l)
• Cell reaction reversed during recharging.
• 2PbSO4(s) 2H2O(l) --gt Pb(s) PbO2(s)
  2H2SO4(aq)

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Lead Storage Battery

• Half-reactions
• anode Pb(s) HSO42-(aq) --gt PbSO4(s) H
  2e-
• cathode PbO2(s) 3H(aq) HSO42-(aq) 2e-
  --gt PbSO4(s) 2H2O(l)
• overall Pb(s) PbO2(s) 2H 2HSO4-(aq) --gt
  2PbSO4(s) 2H2O(l)
• Cell reaction reversed during recharging.

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Lead Storage Battery

• Half-reactions during recharging (nonspontaneous)
• cathode PbSO4(s) H 2e- --gt Pb(s)
  HSO42-(aq)
• anode PbSO4(s) 2H2O(l) --gt
• PbO2(s) 3H(aq) HSO42-(aq) 2e-
• overall 2PbSO4(s) 2H2O(l) --gt
• PbO2(s) Pb(s) 2H 2HSO4-(aq)
• Cell converted into electrolytic cell via
  application of external electrical energy.

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Fuel Cells

• Voltaic-like cell that operates with continuous
  supply of energetic reactants (fuel) to the
  electrodes
• utilize combustion reactions
• do not store chemical energy
• Not self-contained since reactants must be
  supplied to the electrodes
• Example Hydrogen-Oxygen fuel cell

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Hydrogen-Oxygen Fuel Cell
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Hydrogen-Oxygen Fuel Cell

• Half-reactions

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Hydrogen-Oxygen Fuel Cell

• Half-reactions
• anode 2H2(g) 4OH-(aq) --gt 4H2O(l) 4e-

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Hydrogen-Oxygen Fuel Cell

• Half-reactions
• anode 2H2(g) 4OH-(aq) --gt 4H2O(l) 4e-
• cathode O2(g) 2H2O(l) 4e- --gt 4OH-(aq)

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Hydrogen-Oxygen Fuel Cell

• Half-reactions
• anode 2H2(g) 4OH-(aq) --gt 4H2O(l) 4e-
• cathode O2(g) 2H2O(l) 4e- --gt 4OH-(aq)
• overall 2H2(g) O2(g) --gt 2H2O(l)

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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

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

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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 oxidation
• 2Fe(s) 2Na2CrO4(aq) 2H2O(l) --gt
  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

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Standard Electrode Potentials

• Half-reaction E
• F2(g) 2e- -gt 2F-(aq) 2.87 V
• Ag(aq) e- -gt Ag(s) 0.80 V
• Cu2(aq) 2e- -gt Cu(s) 0.34 V
• 2H(aq) 2e- -gt H2(g) 0 V
• Ni2(aq) 2e- -gt Ni(s) -0.25 V
• Fe2(aq) 2e- -gt Fe(s) -0.44 V
• Zn2(aq) 2e- -gt Zn(s) -0.76 V
• Al3(aq) 3e- -gt Al(s) -1.66 V
• Mg2(aq) 2e- -gtMg(s) -2.38 V

Metals more easily oxidized than Fe have more
negative Es
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Cathodic Protection
galvanized steel (Fe)
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Cathodic Protection
(anode)
(cathode)
(electrolyte)
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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 the electrons through the external circuit

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Electrolysis of Molten NaCl
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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
• 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 has passed
• 1 C 1 A x 1 s
• 1 F 1 mole e- 96500 C

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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
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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?

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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.