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Redox potentials

Electrochemistry

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Redox ReactionsOxidation

• loss of electrons

Reduction

• gain of electrons

oxidizing agent

• substance that cause oxidation by being reduced

reducing agent

• substance that cause oxidation by being reduced

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Electrochemistry

In the broadest sense, electrochemistry is the study of chemical reactions that produce electrical effects and of the chemical phenomena that are caused by the action of currents or voltages.

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Voltaic Cells

• harnessed chemical reaction which produces an electric current

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Voltaic Cells

Cells and Cell Reactions

Daniel's Cell

Zn(s) + Cu+2(aq) ---> Zn+2

(aq) + Cu(s)

oxidation half reaction

anode Zn(s) ---> Zn+2(aq) + 2 e-

reduction half reaction

cathode Cu+2(aq) + 2 e- ---> Cu(s)

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Voltaic Cells

• copper electrode dipped into a solution of copper(II) sulfate

• zinc electrode dipped into a solution of zinc sulfate

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Voltaic Cells

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Hydrogen Electrode• consists of a platinum

electrode covered with a fine powder of platinum around which H2(g) is bubbled. Its potential is defined as zero volts.

Hydrogen Half-Cell

H2(g) = 2 H+(aq) + 2 e-

reversible reaction

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Standard Reduction Potentials

• the potential under standard conditions (25oC with all ions at 1 M concentrations and all gases at 1 atm pressure) of a half-reaction in which reduction is occurring

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Some Standard Reduction Potentials Table 18-1, pg 837

Li+ + e- ---> Li -3.045 v

Zn+2 + 2 e- ---> Zn -0.763v

Fe+2 + 2 e- ---> Fe -0.44v

2 H+(aq) + 2 e- ---> H2(g) 0.00v

Cu+2 + 2 e- ---> Cu +0.337v

O2(g) + 4 H+(aq) + 4 e- ---> 2 H2O(l) +1.229v

F2 + 2e- ---> 2 F- +2.87v

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If the reduction of mercury (I) in a voltaic cell is desired, the half reaction is:

Which of the following reactions could be used as the anode (oxidation)?

A, B

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Cell Potential

• the potential difference, in volts, between the electrodes of an electrochemical cell

• Direction of Oxidation-Reduction Reactions

• positive value indicates a spontaneous reaction

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Standard Cell Potential

• the potential difference, in volts, between the electrodes of an electrochemical cell when the all concentrations of all solutes is 1 molar, all the partial pressures of any gases are 1 atm, and the temperature at 25oC

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Cell Diagram

• the shorthand representation of an electrochemical cell showing the two half-cells connected by a salt bridge or porous barrier, such as:

Zn(s)/ZnSO4(aq)//CuSO4(aq)/Cu(s)

anode cathode

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Metal Displacement Reactions

• solid of more reactive metals will displace ions of a less reactive metal from solution

• relative reactivity based on potentials of half reactions

• metals with very different potentials react most vigorously

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Ag+ + e- --->Ag E°= 0.80 V

Cu2+ + 2e- ---> Cu E°= 0.34 V

Will Ag react with Cu2+?

yes, no

Will Cu react with Ag+?

yes, no

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Gibbs Free Energyand Cell Potential

G = - nFE

where n => number of electrons changed

F => Faraday’s constant

E => cell potential

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Applications of Electrochemical Cells

Batteries– device that converts chemical energy into

electricity

Primary Cells– non-reversible electrochemical cell– non-rechargeable cell

Secondary Cells– reversible electrochemical cell– rechargeable cell

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Applications of Electrochemical Cells

Batteries

Primary Cells

"dry" cell & alkaline cell 1.5 v/cell

mercury cell 1.34 v/cell

fuel cell 1.23v/cell

Secondary Cells

lead-acid (automobile battery) 2 v/cell

NiCad 1.25 v/cell

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“Dry” Cell

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“Dry” Cell

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“Flash Light” Batteries

"Dry" CellZn(s) + 2 MnO2(s) + 2 NH4

+ ----->

Zn+2(aq) + 2 MnO(OH)(s) + 2 NH3

Alkaline CellZn(s) + 2 MnO2(s) ---> ZnO(s) + Mn2O3(s)

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“New” Super Iron Battery

Mfe(VI)O4 + 3/2 Zn 1/2 Fe(III)2O3 + 1/2 ZnO + MZnO2

(M = K2 or Ba)Environmentally friendlier than MnO2 containing batteries.

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Lead-Acid(Automobile Battery)

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Lead-Acid(Automobile Battery)

Pb(s) + PbO2(s) + 2 H2SO4 = 2 PbSO4(s) + 2 H2O

2 v/cell

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Nickel-Cadmium (Ni-Cad)

Cd(s) + 2 Ni(OH)3(s) = Cd(OH)2(s) + 2 Ni(OH)2(s)

NiCad 1.25 v/cell

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Automobile Oxygen Sensor

Air, constant [O ]2

porous Pt electrodes

migrating O ions2-

measured potential difference

exhaust gas, unknown [O ]2

ZrO / CaO2

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Automobile Oxygen Sensor

• see Oxygen Sensor Movie from Solid-State Resources CD-ROM

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pH = (Eglass electrode - constant)/0.0592

pH Meter

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Effect of Concentration on Cell Voltage: The Nernst Equation

Ecell = Eocell - (RT/nF)ln Q

Ecell = Eocell - (0.0592/n)log Q

where Q => reaction quotient

Q = [products]/[reactants]

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EXAMPLE: What is the cell potential for the Daniel's cell when the [Zn+2] = 10 [Cu+2] ?

Q = ([Zn+2]/[Cu+2] = (10 [Cu+2])/[Cu+2] = 10

Eo = (0.34 V)Cu couple + (-(-0.76 V)Zn couple

n = 2, 2 electron change Ecell = Eocell - (0.0257/n)ln Q

thus Ecell = (1.10 - (0.0257/2)ln 10) V

Ecell = (1.10 - (0.0257/2)2.303) V

Ecell = (1.10 - 0.0296) V = 1.07 V

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Nernst Equation

+0.83 V

salt bridge

H in2

1 atm

voltmeter

Pt electrode

Pt electrodeNaOH

1 M

H in2

1 atm

e– e–

anode (–)

HCl 1 M

cathode (+)

pn+

+

+

+

–

–

–

–

[H+]acid side [H+]base side

E = Eo – RTnF ln Q =

– 2.3 RTF log

[H+]base side[H+]acid side

[h+]p-type side [h+]n-type side

E (in volts) = – 2.3 RT

F log [h+]n-type side [h+]p-type side