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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.

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Page 1: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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.

Page 2: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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.

Page 3: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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.

Page 4: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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.

Page 5: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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

Page 6: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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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

Page 7: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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.

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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)

Page 9: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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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−

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10

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

Page 11: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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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

Page 12: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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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.

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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-.

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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.

Page 15: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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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.

Page 16: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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16

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 reaction—from

higher to lower

potential energy.

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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

E°cell = Eocathode - E

oanode

Page 18: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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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

Page 19: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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19

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

Page 20: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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20

Electrochemistry

© 2009, Prentice-Hall, Inc.

Standard Cell Potential Pt|H2(g, 1 bar)|H+(a = 1) || Cu2+(1 M)|Cu(s) E°cell = 0.340 V

E°cell = E°cathode - E°anode

E°cell = E°Cu2+/Cu - E°H+/H2

0.340 V = E°Cu2+/Cu - 0 V

E°Cu2+/Cu = +0.340 V

H2(g, 1 atm) + Cu2+(1 M) → H+(1 M) + Cu(s) E°cell = 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

Page 21: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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21

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.

Page 22: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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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

Page 23: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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23

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

Page 24: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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24

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

Page 25: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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25

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

Page 26: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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26

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

Page 27: Chapter 20 Electrochemistry - Yolaholcombslab.yolasite.com/resources/Chapter 20.pdf · Chapter 20 Electrochemistry Chemistry, ... Chapter 20 Problems • Problems 15, 17, 19, 23,

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27

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.

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Electrochemistry

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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

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Batteries

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Electrochemistry

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Alkaline Batteries

Electrochemistry

© 2009, Prentice-Hall, Inc.

Hydrogen Fuel Cells

Electrochemistry

© 2009, Prentice-Hall, Inc.

Corrosion and…

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Electrochemistry

© 2009, Prentice-Hall, Inc.

…Corrosion Prevention