Chapter 20 Electrochemistry - 20.pdfChapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23, 27, ... 5 Electrochemistry

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  • 7/3/2012

    1

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Chapter 20

    Electrochemistry

    Chemistry, The Central Science, 11th edition

    Theodore L. Brown; H. Eugene LeMay, Jr.;

    and Bruce E. Bursten

    John D. Bookstaver

    St. Charles Community College

    Cottleville, MO

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Chapter 20 Problems

    Problems 15, 17, 19, 23, 27, 29, 33, 39,

    59

    Electrochemistry

    2009, Prentice-Hall, Inc.

    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.

  • 7/3/2012

    2

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Oxidation-Reduction Reactions

    Basic Terms

    Oxidation - loss of electrons

    Reduction - gain of electrons

    Redox reaction

    oxidizing agent - substance that causes

    oxidation by being reduced

    reducing agent - substance that causes

    reduction by being oxidized

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Electrochemical Reactions

    In electrochemical reactions, electrons

    are transferred from one species to

    another.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Oxidation Numbers (State)

    In order to keep

    track of what loses

    electrons and what

    gains them, we

    assign oxidation

    numbers.

  • 7/3/2012

    3

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Oxidation and Reduction

    A species is oxidized when it loses electrons.

    Here, zinc loses two electrons to go from neutral

    zinc metal to the Zn2+ ion.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Oxidation and Reduction

    A species is reduced when it gains electrons.

    Here, each of the H+ gains an electron, and they

    combine to form H2.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Oxidation and Reduction

    What is reduced is the oxidizing agent.

    H+ oxidizes Zn by taking electrons from it.

    What is oxidized is the reducing agent.

    Zn reduces H+ by giving it electrons.

  • 7/3/2012

    4

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Oxidation and Reduction

    Oxidation

    O.S. of some element increases in the

    reaction.

    Electrons are on the right of the equation

    Reduction

    O.S. of some element decreases in the

    reaction.

    Electrons are on the left of the equation.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Assigning Oxidation Numbers

    1. Elements in their elemental form have

    an oxidation number of 0.

    2. The oxidation number of a monatomic

    ion is the same as its charge.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Assigning Oxidation Numbers

    3. Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.

    Oxygen has an oxidation number of 2, except in the peroxide ion, which has an oxidation number of 1.

    Hydrogen is 1 when bonded to a metal and +1 when bonded to a nonmetal.

  • 7/3/2012

    5

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Assigning Oxidation Numbers

    3. Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.

    Fluorine always has an oxidation number of 1.

    The other halogens have an oxidation number of 1 when they are negative; they can have positive oxidation numbers, however, most notably in oxyanions.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Assigning Oxidation Numbers

    4. The sum of the oxidation numbers in a

    neutral compound is 0.

    5. The sum of the oxidation numbers in a

    polyatomic ion is the charge on the

    ion.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Oxidation State

    What is the oxidation state of S in H2SO4?

    H => +1

    O => -2

    neutral compound, thus sum equals

    zero

    4O => 4*-2 = -8

    2H => 2*+1 = +2

    0 = +2 + (x) + (-8)

    x = +6

  • 7/3/2012

    6

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Oxidation State

    What is the oxidation state of Cl in

    HClO4?

    H => +1

    O => -2

    neutral compound, thus sum equals

    zero

    4O => 4*-2 = -8

    H => 1*+1 = +1

    0 = +1 + (y) + (-8)

    y = +7

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Oxidation and Reduction Half-

    Reactions A reaction represented by two half-

    reactions.

    Oxidation:

    Reduction:

    Overall:

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

    Cu2+(aq) + 2 e- Cu(s)

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

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Balancing Oxidation-Reduction

    Equations

    Few can be balanced by inspection.

    Systematic approach required.

    The Half-Reaction (Ion-Electron)

    Method

  • 7/3/2012

    7

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Balancing Oxidation-Reduction

    Equations

    Perhaps the easiest way to balance the

    equation of an oxidation-reduction

    reaction is via the half-reaction method.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Balancing Oxidation-Reduction

    Equations

    This involves treating (on paper only) the

    oxidation and reduction as two separate

    processes, balancing these half reactions,

    and then combining them to attain the

    balanced equation for the overall reaction.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    The Half-Reaction Method

    1. Assign oxidation numbers to

    determine what is oxidized and what is

    reduced.

    2. Write the oxidation and reduction half-

    reactions.

  • 7/3/2012

    8

    Electrochemistry

    2009, Prentice-Hall, Inc.

    The Half-Reaction Method

    3. Balance each half-reaction.

    a. Balance elements other than H and O.

    b. Balance O by adding H2O.

    c. Balance H by adding H+.

    d. Balance charge by adding electrons.

    4. Multiply the half-reactions by integers so that the electrons gained and lost are the same.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    The Half-Reaction Method

    5. Add the half-reactions, subtracting things that appear on both sides.

    6. Make sure the equation is balanced according to mass.

    7. Make sure the equation is balanced according to charge.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    The Half-Reaction Method

    Consider the reaction between MnO4 and C2O4

    2 :

    MnO4

    (aq) + C2O42

    (aq) Mn2+

    (aq) + CO2 (aq)

  • 7/3/2012

    9

    Electrochemistry

    2009, Prentice-Hall, Inc.

    The Half-Reaction Method

    First, we assign oxidation numbers.

    MnO4 + C2O4

    2- Mn2+ + CO2

    +7 +3 +4 +2

    Since the manganese goes from +7 to +2, it is reduced.

    Since the carbon goes from +3 to +4, it is oxidized.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Oxidation Half-Reaction

    C2O42 CO2

    To balance the carbon, we add a

    coefficient of 2:

    C2O42 2 CO2

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Oxidation Half-Reaction

    C2O42 2 CO2

    The oxygen is now balanced as well. To

    balance the charge, we must add 2

    electrons to the right side.

    C2O42 2 CO2 + 2 e

  • 7/3/2012

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    Electrochemistry

    2009, Prentice-Hall, Inc.

    Reduction Half-Reaction

    MnO4 Mn2+

    The manganese is balanced; to balance

    the oxygen, we must add 4 waters to

    the right side.

    MnO4 Mn2+ + 4 H2O

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Reduction Half-Reaction

    MnO4 Mn2+ + 4 H2O

    To balance the hydrogen, we add 8 H+

    to the left side.

    8 H+ + MnO4 Mn2+ + 4 H2O

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Reduction Half-Reaction

    8 H+ + MnO4 Mn2+ + 4 H2O

    To balance the charge, we add 5 e to

    the left side.

    5 e + 8 H+ + MnO4 Mn2+ + 4 H2O

  • 7/3/2012

    11

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Combining the Half-Reactions

    Now we evaluate the two half-reactions

    together:

    C2O42 2 CO2 + 2 e

    5 e + 8 H+ + MnO4 Mn2+ + 4 H2O

    To attain the same number of electrons

    on each side, we will multiply the first

    reaction by 5 and the second by 2.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Combining the Half-Reactions

    5 C2O42 10 CO2 + 10 e

    10 e + 16 H+ + 2 MnO4 2 Mn2+ + 8 H2O

    When we add these together, we get:

    10 e + 16 H+ + 2 MnO4 + 5 C2O4

    2

    2 Mn2+ + 8 H2O + 10 CO2 +10 e

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Combining the Half-Reactions

    10 e + 16 H+ + 2 MnO4 + 5 C2O4

    2

    2 Mn2+ + 8 H2O + 10 CO2 +10 e

    The only thing that appears on both sides are the electrons. Subtracting them, we are left with:

    16 H+ + 2 MnO4 + 5 C2O4

    2

    2 Mn2+ + 8 H2O + 10 CO2

  • 7/3/2012

    12

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Balancing in Basic Solution

    If a reaction occurs in basic solution, one

    can balance it as if it occurred in acid.

    Once the equation is balanced, add OH

    to each side to neutralize the H+ in the

    equation and create water in its place.

    If this produces water on both sides, you

    might have to subtract water from each

    side.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Important Electrochemical Terms

    An electrochemical cell is a device that combines

    two half-cells with the appropriate connections

    between electrodes and solutions

    A voltaic cell is an electrochemical cell in which

    electric current is generated from a spontaneous

    redox reaction

    The anode is the electrode at which oxidation occurs

    and the cathode is the electrode at which reduction

    occurs

    EOS

    The cell potential (Ecell) is the potential difference

    that propels electrons from the anode to the cathode

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Voltaic Cells

    In spontaneous

    oxidation-reduction

    (redox) reactions,

    electrons are

    transferred and

    energy is released.

  • 7/3/2012

    13

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Cell Diagrams Conventions

    Place the anode on the left side of the diagram

    Place the cathode on the right side of the diagram

    Use a single vertical line ( | ) to represent the

    boundary between different phases, such as

    between an electrode and a solution

    EOS

    Use a double vertical line ( || ) to represent a salt

    bridge or porous barrier separating two half-

    cells

    Electrochemistry

    2009, Prentice-Hall, Inc.

    An Example Cell Diagram

    EOS

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Terminology

    Galvanic (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-.

  • 7/3/2012

    14

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Voltaic Cells

    harnessed chemical reaction which

    produces an electric current

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Terminology

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

    Ecell = 1.103 V

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Voltaic Cells

    We can use that

    energy to do work if

    we make the

    electrons flow

    through an external

    device.

    We call such a

    setup a voltaic cell.

  • 7/3/2012

    15

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Voltaic Cells

    A typical cell looks

    like this.

    The oxidation

    occurs at the anode.

    The reduction

    occurs at the

    cathode.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Voltaic Cells

    Once even one

    electron flows from

    the anode to the

    cathode, the

    charges in each

    beaker would not be

    balanced and the

    flow of electrons

    would stop.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Voltaic Cells

    Therefore, we use a

    salt bridge, usually a

    U-shaped tube that

    contains a salt

    solution, to keep the

    charges balanced.

    Cations move toward

    the cathode.

    Anions move toward

    the anode.

  • 7/3/2012

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    Electrochemistry

    2009, Prentice-Hall, Inc.

    Voltaic Cells In the cell, then,

    electrons leave the

    anode and flow

    through the wire to

    the cathode.

    As the electrons

    leave the anode,

    the cations formed

    dissolve into the

    solution in the

    anode

    compartment.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Voltaic Cells As the electrons

    reach the cathode,

    cations in the

    cathode are

    attracted to the now

    negative cathode.

    The electrons are

    taken by the cation,

    and the neutral

    metal is deposited

    on the cathode.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Electromotive Force (emf)

    Water only

    spontaneously flows

    one way in a

    waterfall.

    Likewise, electrons

    only spontaneously

    flow one way in a

    redox reactionfrom

    higher to lower

    potential energy.

  • 7/3/2012

    17

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Electromotive Force (emf)

    The potential difference between the

    anode and cathode in a cell is called the

    electromotive force (emf).

    It is also called the cell potential and is

    designated Ecell.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Cell Potential

    Cell potential is measured in volts (V).

    1 V = 1 J

    C

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Cell Potential (Ecell)

    The cell potential (Ecell) is the potential

    difference that propels electrons from

    the anode to the cathode

    Ecell = Eocathode - Eo

    anode

  • 7/3/2012

    18

    Electrochemistry

    2009, Prentice-Hall, Inc.

    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 for the Standard Hydrogen Electrode (SHE)

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Standard Hydrogen Electrode

    Half-cell values are referenced to a standard

    hydrogen electrode (SHE).

    By definition, the reduction potential for

    hydrogen is 0 V:

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

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Standard Hydrogen Electrode

    In the standard

    hydrogen electrode,

    hydrogen gas at exactly

    1 bar pressure is

    bubbled over an inert

    platinum electrode and

    into an aqueous solution

    with the concentration

    adjusted so that the

    activity of H3O+ is

    exactly equal to one EOS

  • 7/3/2012

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    Electrochemistry

    2009, Prentice-Hall, Inc.

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

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

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Standard Electrode Potential, E

    E defined by international agreement.

    The tendency for a reduction process 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).

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Standard Electrode Potentials

    A standard electrode potential, Eo, is based on

    the tendency for reduction to occur at the electrode

    The cell voltage, called the standard cell potential

    (Eocell), is the difference between the standard

    potential of the cathode and that of the anode

    EOS

    Two illustrations Cu and Zn

    Voltaic cells can produce electrical work

    w = n F Ecell

    VideoClip

    ppt_Heitz/AABTHAZ0.MOV

  • 7/3/2012

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    Electrochemistry

    2009, Prentice-Hall, Inc.

    Standard Cell Potential 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

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Measuring Standard Reduction

    Potential

    anode anode cathode cathode

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Cu2+/Cu Electrode

    EOS

  • 7/3/2012

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    Electrochemistry

    2009, Prentice-Hall, Inc.

    Zn2+/Zn Electrode

    EOS

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Observed Voltages

    EOS

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Standard Reduction Potentials

    Reduction

    potentials for

    many

    electrodes

    have been

    measured and

    tabulated.

  • 7/3/2012

    22

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Important Points About

    Electrode and Cell Potentials

    Electrode potentials and cell voltages are intensive

    properties independent of the amount of matter

    Cell voltages can be ascribed to oxidation

    reduction reactions without regard to voltaic cells

    EOS

    Ecell = Ecathode Eanode

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Cell Potentials

    For the oxidation in this cell,

    For the reduction,

    Ered = 0.76 V

    Ered = +0.34 V

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Cell Potentials

    Ecell = Ered (cathode) Ered (anode)

    = +0.34 V (0.76 V)

    = +1.10 V

  • 7/3/2012

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    Electrochemistry

    2009, Prentice-Hall, Inc.

    Oxidizing and Reducing Agents

    The strongest

    oxidizers have the

    most positive

    reduction potentials.

    The strongest

    reducers have the

    most negative

    reduction potentials.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Oxidizing and Reducing Agents

    The greater the

    difference between

    the two, the greater

    the voltage of the

    cell.

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Criteria for Spontaneous

    Change

    If Ecell is positive, the reaction in the forward

    direction (from left to right) is spontaneous

    If Ecell is negative, the reaction is nonspontaneous

    If Ecell = 0, the system is at equilibrium

    EOS

    When a cell reaction is reversed, Ecell and DG

    change signs

  • 7/3/2012

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    Electrochemistry

    2009, Prentice-Hall, Inc.

    Free Energy

    DG for a redox reaction can be found by

    using the equation

    DG = nFE

    where n is the number of moles of

    electrons transferred, and F is a

    constant, the Faraday.

    1 F = 96,485 C/mol = 96,485 J/V-mol

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Free Energy

    Under standard conditions,

    DG = nFE

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Nernst Equation

    Remember that

    DG = DG + RT ln Q

    This means

    nFE = nFE + RT ln Q

  • 7/3/2012

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    Electrochemistry

    2009, Prentice-Hall, Inc.

    Nernst Equation

    Dividing both sides by nF, we get the

    Nernst equation:

    E = E RT

    nF ln Q

    or, using base-10 logarithms,

    E = E 2.303 RT

    nF log Q

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Nernst Equation

    At room temperature (298 K),

    Thus the equation becomes

    E = E 0.0592

    n log Q

    2.303 RT

    F = 0.0592 V

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Summary of Important

    Relationships

    EOS

  • 7/3/2012

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    Electrochemistry

    2009, Prentice-Hall, Inc.

    Concentration and Cell

    Voltage If the discussion is limited to 25 oC, the equation is

    o ocell cell cell

    0.0592ln ln

    RTE E Q E Q

    nF n

    The Nernst equation relates a cell voltage for

    nonstandard conditions (Ecell) to a standard cell voltage,

    Eocell, and the concentrations of reactants and products

    EOS

    The Nernst equation is useful for determining the

    concentration of a species in a voltaic cell through

    a measurement of Ecell

    Electrochemistry

    2009, Prentice-Hall, Inc.

    A Sample Problem

    ocell cell

    0.0592ln E E Q

    n

    EOS

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Pt|Fe2+(0.10 M),Fe3+(0.20 M)||Ag+(1.0 M)|Ag(s)

    Applying the Nernst Equation for Determining Ecell. What

    is the value of Ecell for the voltaic cell pictured below and

    diagrammed as follows?

    EXAMPLE

  • 7/3/2012

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    Electrochemistry

    2009, Prentice-Hall, Inc.

    Ecell = Ecell - log Q n

    0.0592 V

    Pt|Fe2+(0.10 M),Fe3+(0.20 M)||Ag+(1.0 M)|Ag(s)

    Ecell = Ecell - log n

    0.0592 V [Fe3+]

    [Fe2+] [Ag+]

    Fe2+(aq) + Ag+(aq) Fe3+(aq) + Ag (s)

    Ecell = 0.029 V 0.018 V = 0.011 V

    EXAMPLE

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Concentration Cells

    If the cell potential is

    determined solely by a

    difference in the

    concentration of solutes

    in equilibrium with

    identical electrodes, that

    cell is called a

    concentration cell

    EOS

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Concentration Cells

    Notice that the Nernst equation implies that a cell

    could be created that has the same substance at

    both electrodes.

    For such a cell, would be 0, but Q would not. Ecell

    Therefore, as long as the concentrations

    are different, E will not be 0.

  • 7/3/2012

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    Electrochemistry

    2009, Prentice-Hall, Inc.

    pH Measurement

    The Nernst equation can be rewritten in terms of

    concentration cells using hydrogen electrodes

    EOS

    Ecell = (0.0592)(pH)

    One useful concentration cell is the hydrogen cell

    it is directly related to pH

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Applications of

    Oxidation-Reduction

    Reactions

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Batteries

  • 7/3/2012

    29

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Alkaline Batteries

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Hydrogen Fuel Cells

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Corrosion and

  • 7/3/2012

    30

    Electrochemistry

    2009, Prentice-Hall, Inc.

    Corrosion Prevention

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