chapter 20 petrucci

5/26/2018 Chapter 20 Petrucci

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Slide 1 of 53 Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 20Slide 1 of 53

PHILIP DUTTONUNIVERSITY OF WINDSOR

DEPARTMENT OF CHEMISTRY ANDBIOCHEMISTRY

TENTH EDITION

GENERAL CHEMISTRYPrinciples and Modern Applications

PETRUCCI HERRING MADURA BISSONNETTE

Electrochemistry 20


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Slide 2 of 53

Spontaneous Change:

Entropy and Gibbs Energy

Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 20Slide 2 of 53

ONTENTS19-1 Electrode Potentials and Their

Measurement

19-2 Standard Electrode Potentials

19-3 Ecell, G, and K

19-4 Ecellas a function of

Concentrations

19-5 Batteries: Producing ElectricityThrough Chemical Reactions

19-6 Corrosion: Unwanted Voltaic

Cells

19-7 Electrolysis: Causing

Nonspontaneous Reactions toOccur

19-8 Industrial Electrolysis Processes


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Slide 3 of 53

20-1 Electrode Potentials and Their Measurement


FIGURE 20-1

Behaviour of Ag+(aq) and Zn+(aq) in the presence of copper

Cu(s) + 2Ag+(aq)

Cu2+(aq) + 2 Ag(s)

Cu(s) + Zn2+(aq)

No reaction


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Slide 4 of 53

An electrochemical half cell

FIGURE 20-2

Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 20

Anode

Cathode


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Slide 5 of 53

Measurement of the electromotive force of an electrochemical cell

FIGURE 20-3



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The reaction Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s) in an electrochemical cell

FIGURE 20-4


Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

Ecell= 1.103 V


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Slide 7 of 53

Cell Diagrams and Terminology


Electromotive force, Ecell

The cell voltage or cell potential.

Cell diagram

Shows the components of the cell in a symbolic way.

Anode (where oxidation occurs) on the left.

Cathode (where reduction occurs) on the right.

Boundary between phases shown by |.

Boundary between half cells(commonly a salt bridge) shown by ||.


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The reaction Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s) in an electrochemical cell

FIGURE 20-4


Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)

Ecell= 1.103 V


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Slide 9 of 53

Galvanic (or voltaic) cells

Produce electricity as a result of spontaneous reactions.

Electrolytic cells

Non-spontaneous chemical change driven by electricity.

Couple, M|Mn+

A pair of species related by a change in number of e-.



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Slide 12 of 53

20-2 Standard Electrode Potentials

Cell voltages, the potential differences

between electrodes, are among the most

precise scientific measurements.

The potential of an individual electrode is

difficult to establish.

Arbitrary zero is chosen.


The Standard Hydrogen Electrode (SHE)


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Slide 13 of 53

The standard hydrogen electrode (SHE)

FIGURE 20-5


2 H+(a = 1) + 2 e- H2(g, 1 bar) E= 0

V

Pt|H2(g, 1 bar)|H+(a = 1)


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Slide 14 of 53

Standard Electrode Potential, E

Ecell= Ecathode(right)Eanode,(left)

The tendency for a reductionprocess to occur at an electrode.

All ionic species present at a=1 (approximately 1 M).

All gases are at 1 bar (approximately 1 atm).

Where no metallic substance is indicated, the potential is

established on an inert metallic electrode (ex. Pt).



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Slide 15 of 53 Copyright 2011 Pearson Canada Inc.General Chemistry: Chapter 20

Cu2+(1M) + 2 e- Cu(s) ECu2+/Cu= ?

Pt|H2(g, 1 bar)|H+(a = 1) || Cu2+(1 M)|Cu(s) Ecell= 0.340 V

Standard cell potential: the potential difference of a

cell formed from twostandardelectrodes.

E

cell=E

cathode -E

anode

cathodeanode


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Pt|H2(g, 1 bar)|H+(a = 1) || Cu2+(1 M)|Cu(s) Ecell= 0.340 V

Ecell=Ecathode -Eanode

Ecell=ECu2+/Cu -EH+/H2

0.340 V =ECu2+/Cu -0 V

ECu2+/Cu = +0.340 V

H2(g, 1 atm) + Cu2+(1 M) H+(1 M) + Cu(s) Ecell= 0.340

V


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Slide 17 of 53

Measuring standard reduction potential

FIGURE 20-6


anodeanode cathode cathode


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Slide 18 of 53

TABLE 20.1 Some Selected Standard Electrode (Reduction)

Potentials at 25C


Reduction Half-Reaction E, V

Acidic Solution


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Slide 19 of 53


Potentials at 25C (continued)



Acidic Solution


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Slide 20 of 53


Potentials at 25C (continued)



Acidic Solution

Basic Solution


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Slide 23 of 53

20-3 Ecell, G, and Keq

Faraday constant,

F = 96,485 C mol-1

When products are in their

standard states


elec= -zFEcellG= -zFEcell

G= -zFEcell

Michael Faraday 1791-1867


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Slide 24 of 53

Combining Reduction Half-Equations


Fe3+(aq) + 3e- Fe(s) EFe3+/Fe= ?

Fe2+(aq) + 2e- Fe(s) EFe2+/Fe= -0.440 V

Fe3+

(aq) + 1e-

Fe2+

(aq) E

Fe3+/Fe2+= 0.771 V

Fe3+(aq) + 3e- Fe(s)

G= +0.880 J

G

= -0.771 J

G= +0.109 JEFe3+/Fe= +0.331 V

G= +0.109 J = -nFE

EFe3+/Fe= +0.109 J /(-3F)= -0.0363 V

but cannot simply add E

can add G


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Slide 25 of 53

Spontaneous Change in

Oxidation-Reduction Reactions


G < 0 for spontaneous change.

ThereforeEcell> 0 because Gcell= -nFEcell

E cell> 0Reaction proceeds spontaneously as written.

E cell= 0Reaction is at equilibrium.

E cell< 0Reaction proceeds in the reversedirection spontaneously.


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Slide 28 of 53

The Behavior or Metals Toward Acids


M(s) M2+(aq) + 2 e- E= -EM2+/M

2 H+(aq) + 2 e- H2(g) EH+/H2= 0 V

2 H+(aq) + M(s) H2(g) + M2+(aq)

Ecell=EH+/H2-EM2+/M= -EM2+/M

WhenEM2+/M < 0,Ecell> 0. Therefore G< 0.

Metals with negative reduction potentials react with acids.


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Relationship Between Ecelland Keq


G= -RTlnKeq= -zFEcell

Ecell = zF

RT

lnKeq

Ecell =z

0.25693lnKeq


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Slide 31 of 53

A summary of important thermodynamic, equilibrium and electrochemical

relationships under standard conditions.

FIGURE 20-8



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Slide 33 of 53

20-4 Ecellas a Function of Concentration

Copyright 2011 Pearson Canada Inc.

General Chemistry: Chapter 20

FIGURE 20-9

Variation of Ecellwith ion concentrations

G= G-RTln Q

-zFEcell= -zFEcell-RTln Q

Ecell=Ecell- ln QzF

RT

Convert to log10and calculate constants.

Ecell=Ecell- log Qz

0.0592 VThe Nernst Equation


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Slide 34 of 53

Concentration Cells



FIGURE 20-11

A concentration cell

Two half cells with identical electrodes but different ion concentrations.

2 H+(1 M) 2 H+(xM)

Pt|H2(1 atm)|H+(xM)||H+(1.0 M)|H2(1 atm)|Pt(s)

2 H+(1 M) + 2 e- H2(g, 1 atm)

H2(g, 1 atm) 2 H+(xM) + 2 e-

Ecell=EH+/H2-EH+/H2

= 0


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Slide 35 of 53Copyright 2011 Pearson Canada Inc.

General Chemistry: Chapter 20Slide 35 of 54

Ecell=Ecell- logz

0.0592 V x2

12

Ecell= 0 - log2

0.0592 V x2

1

Ecell= - 0.0592 V logx

Ecell= (0.0592 V) pH

2 H+(1 M) 2 H+(xM)

Ecell=Ecell

- log Qz

0.0592 V

Ecell=EH+/H2-EH+/H2

= 0

but we can calculateusing the Nernst Equation


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Slide 38 of 53

Measurement of Ksp



FIGURE 20-12

A concentration cell for determining Kspof AgI

Ag+(0.100 M) Ag+(satd M)

Ag|Ag+(satd AgI)||Ag+(0.10 M)|Ag(s)

Ag+(0.100 M) + e- Ag(s)

Ag(s) Ag+(satd) + e-

Work Example 20-11 as an exerciseto understand the process.


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Schematic diagrams of some common electrodes

FIGURE 20-13



0.22233 V

0.2680 V (satd KCl)

or

0.2412 V (1 M KCl)

Th Gl El t d d th El t h i l


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Slide 41 of 53

The Glass Electrode and the Electrochemical

Measurement of pH



Ag|AgCl(s)|Cl-(1.0M),H+(1.0M)|glass membrane|H+(unknown)|| Cl-(1.0 M)|AgCl(s)|Ag(s)

Ag(s) + Cl- AgCl(s) + e-

H+(1.0 M) H+(unknown)

AgCl(s) + e- Ag(s) + Cl-(aq)

G = G(unknown) G(1.0M)

= G+ RTln[unknown]G- RTln(1.0)

=RTln[unknown]

Ecell= -RTln[unknown]/zF

pH = -log[unknown]=zFEcell/RT


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Slide 42 of 53

20-5 Batteries: Producing Electricity Through

Chemical Reactions

Primary Cells (or batteries).

Cell reaction is not reversible.

Secondary Cells.Cell reaction can be reversed by passing

electricity through the cell (charging).

Flow Batteries and Fuel Cells.

Materials pass through the battery which converts

chemical energy to electric energy.




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The Leclanch (dry) cell

FIGURE 20-14




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The Leclanch Dry Cell


Zn(s) Zn2+(aq) + 2 e-Oxidation:

2 MnO2(s) + H2O(l) + 2 e- Mn2O3(s) + 2 OH

-Reduction:

NH4+ + OH- NH3(g) + H2O(l)Acid-base reaction:

NH3 + Zn2+(aq)+ Cl- [Zn(NH3)2]Cl2(s)

Precipitation reaction:


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Zn2+(aq)+ 2 OH- Zn (OH)2(s)

Zn(s) Zn2+(aq) + 2 e-

Oxidation reaction can be thought of in two steps:

2 MnO2(s) + H2O(l) + 2 e- Mn2O3(s) + 2 OH-Reduction:

Zn(s)+ 2 OH- Zn (OH)2(s) + 2 e-

The alkaline cell


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The Lead-Acid(Storage) Battery


FIGURE 20

The lead-acid (storage) cell


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PbO2(s) + 3 H+(aq) + HSO4

-(aq) + 2 e- PbSO4(s) + 2 H2O(l)

Oxidation:

Reduction:

Pb (s) + HSO4-(aq) PbSO4(s) + H

+(aq) + 2 e-

PbO2(s) + Pb(s) + 2 H+(aq) + HSO4

-(aq) 2 PbSO4(s) + 2 H2O(l)

Ecell=EPbO2/PbSO4-EPbSO4/Pb= 1.74 V(-0.28 V) =

2.02 V


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Slide 48 of 53

The Silver-Zinc Cell: A Button Battery


FIGURE 20-16

The silver-zinc button (miniature) cell

Zn(s) + Ag2O(s) ZnO(s) + 2 Ag(s) Ecell= 1.8 V

Zn(s),ZnO(s)|KOH(satd)|Ag2O(s),Ag(s)

The Nickel Cadmium Cell: A Rechargeable


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The Nickel-Cadmium Cell: A Rechargeable

Battery


Cd(s) + 2 NiO(OH)(s) + 2 H2O(L) 2 Ni(OH)2(s) + Cd(OH)2(s)

A rechargeable nickel-cadmium cell, or nicad battery


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The Lithium-Ion Battery


FIGURE 20-17

The electrodes of a lithium-ion battery


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The positive electrode consists of lithium cobalt(III) oxide, LiCoO2, and the

negative electrode is highly crystallized graphite. To complete the battery an

electrolyte is needed, which can consist of an organic solvent and ions, such as

LiPF6. The structure of LiCoO2, and graphite electrode is illustrated in Figure 20-17.

In the charging cycle at the positive electrode, lithium ions are released into

the electrolyte solution as electrons are removed from the electrode. To

maintain a charge balance, one cobalt(III) ion is oxidized to cobalt(IV) foreach lithium ion released.

LiCoO2(s)+Li(1-x) = 2CoO2(s) + xLi+(solvent) + x e-

C(s) + xLi+(solvent) + x e- = LixC(S)

The layered graphite electrode is shown with lithium ions (violet) intercalated. The

LiCoO2is shown as a face-centered cubic lattice, with the oxygen atoms (red)

occupying the corners and the faces, the cobalt atoms (pink) occupying half of the

edges, and the lithium atoms occupying half of the edges and the central octahedral

hole. This arrangement leads to planes of oxygen, cobalt, oxygen, lithium, oxygen,

cobalt, and oxygen atoms, as indicated in the figure.


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Fuel Cells


FIGURE 20-18

A hydrogen-oxygen fuel cell

O2(g) + 2 H2O(l) + 4 e-

4 OH-

(aq)

2{H2(g) + 2 OH-(aq) 2 H2O(l) + 2 e

-}

2H2(g) + O2(g) 2 H2O(l)

Ecell=EO2/OH--EH2O/H2

= 0.401 V(-0.828 V) = 1.229 V

= G/ H=

0.83


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Slide 53 of 53

Air Batteries


FIGURE 20-19

A simplified aluminum-air battery

4 Al(s) + 3 O2(g) + 6 H2O(l) + 4 OH- 4 [Al(OH)4](aq)


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20-6 Corrosion: Unwanted Voltaic Cells


O2(g) + 2 H2O(l) + 4 e- 4 OH-(aq)

2 Fe(s) 2 Fe2+(aq) + 4 e-

2 Fe(s) + O2(g)+ 2 H2O(l)

2 Fe2+(aq) + 4 OH-(aq)

Ecell= 0.841 V

EO2/OH-= 0.401 V

EFe/Fe2+= -0.440 V

I n neutral solution:

I n acidic solution:

O2(g) + 4 H+(aq) + 4 e- 4 H2O (aq) EO2/OH-= 1.229 V


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Demonstration of corrosion and methods of corrosion protection

FIGURE 20-20


The pink color

results from the

indicatorphenolphthalein in

the presence of base.

The dark blue color

results from the

formation of

Turnbulls blue

KFe[Fe(CN)6].

Zn

Cu


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Protection of iron against electrolytic corrosion

FIGURE 20-21



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The small cylindrical bars of magnesium attached to the steel ship

provide cathodic protection against corrosion.

Magnesium sacrificial anodes


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20-7 Electrolysis: Causing Non-spontaneous

Reactions to Occur


Voltaic Cell :

Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s) EO2/OH-= 1.103 V

Electolytic Cell:

Cu(s) + Zn2+(aq) Cu2+(aq) + Zn(s) EO2/OH-= -1.103 V


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Predicting Electrolysis Reaction


FIGURE 20-22

An electrolytic cell

An Electrolytic Cell

e- is the reverse of the

voltaic cell.

The battery must have a

voltage in excess of 1.103V in order to force the

non-spontaneous reaction.


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Complications in Electrolytic Cells

Overpotential.

Competing reactions.

Non-standard states.

Nature of electrodes.


Q f


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Quantitative Aspects of Electrolysis


ne-= It

F

1 mol e-= 96485 C

Charge (C) = current (C/s) time (s)


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20 8 I d t i l El t l i P


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20-8 Industrial Electrolysis Processes


The refining of copper by electrolysis.


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Electrorefining


Electroplating

Electrosynthesis

A rack of metal parts being lifted from theelectrolyte solution after electroplating.

Chl Alk li P


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Chlor-Alkali Process


FIGURE 20-24

A diaphragm chlor-alkali cell

FIGURE 20-25

The mercury-cell chlor-alkali process

E d f Ch t Q ti


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End of Chapter Questions

Dont just read examples, workthem!!

If you write:

Information is going through your fingers,

Your muscles,

Your nerves,

Directlyto your brain.

Physically experience the solution.

Your eyes and ears are not enough.

chapter 20 petrucci

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