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18Introduction to Electrochemistry

CHAPTER

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Oxidation Oxidation ：氧化反應 ：氧化反應 ReductionReduction ：還原反應：還原反應 Reducing agent : Reducing agent : 還原劑 還原劑 還原其他物種，自身進行氧化反應還原其他物種，自身進行氧化反應 Oxidizing agent Oxidizing agent ：氧化劑：氧化劑 氧化其他物種，自身進行還原反應氧化其他物種，自身進行還原反應

Example 18-1The following reactions are spontaneous and thus proceed to the right, as writtenWhat can we deduce regarding the strengths of H+, Ag+, Cd2+, Zn2+ as electron acceptors? (or oxidizing agents)

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2H+ + Cd(s) H2 + Cd2+

2Ag+ + H2(g) 2Ag(s) + 2H+

Cd2+ + Zn(s) Cd(s) + Zn2+

Ag+ > H+ > Cd2+ > Zn2+

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Figure 18-1Photograph of a “ silver tree.”

Ag+ + e Ag(s)

銅片浸入硝酸銀水溶液中

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Galvanic Cell 賈法尼電池

19.2

spontaneousredox reaction

anodeoxidation

cathodereduction

Chapter18 p495Figure 18-2(a)A galvanic cell at open circuit;

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(b) a galvanic cell doing work;

Chapter18 p495(c) an electronlytic cell. 電解電池

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電池表示方法電池表示方法18B-3 Representing Cells Schematically

Chemists frequently use a shorthand notation to describe electrochemical cells. The cell in Figure 18-2a, for example, is described by

single vertical line indicates a phase boundary, or interface, at which a potential develops.

The double vertical line represents two phase boundaries, one at each end of the salt bridge. A liquid-junction potential develops at each of these interfaces.

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陰極 (cathode)

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Figure 18-3Movement of charge in a galvanic cell.

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18C Electrode Potentials 電極 電位

The cell potential Ecell is related to the free energy of the reaction ΔG by

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If the reactants and products are in their standard states, the resulting cell potential is called the standard cell potential.

where R is the gas constant and T is the absolute temperature.

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(a)

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(b)

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(c)

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Figure 18-5Cell potential in the galvanic cell of Figure 18-4b as a function of time. The cell current, which is directly related to the cell potential, also decreases with the same time behavior.

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If we always follow this convention, the value of Ecell is a measure of the tendency of the cell reaction to occur spontaneously in the direction written from left to right.

the spontaneous cell reaction will occur.

we may write the cell potential Ecell as

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18C-2 The Standard Hydrogen Reference Electrode

an electrode must be easy to construct, reversible, and highly reproducible in its behavior. The standard hydrogen electrode (SHE) meets these specifications and has been used throughout the world for many years as a universal reference electrode. It is a typical gas electrode.

The half-reaction responsible for the potential that develops at this electrode is

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Figure 18-6The hydrogen gas electrode.

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By convention, the potential of the standard hydrogen electrode is assigned a value of 0.000 V at all temperatures.

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Standard Electrode Potentials 標準電極電位

19.3

Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)

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

Zn (s) Zn2+ (1 M) + 2e-Anode (oxidation):

Cathode (reduction):

Zn (s) + 2H+ (1 M) Zn2+ + H2 (1 atm)

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18C-3 Electrode Potential and Standard Electrode Potential

An electrode potential is defined as the potential of a cell in which the electrode in question is the right-hand electrode and the standard hydrogen electrode is the left-hand electrode.

The cell potential is

EAg is the potential of the silver electrode.

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The standard electrode potential, E0, of a half-reaction is defined as its electrode potential when the activities of the reactants and products are all unity.

the E0 value for the half-reaction

the cell shown in Figure 18-7 can be represented schematically as

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Figure 18-7Measurement of the electrode potential for an Ag electrode. If the silver ion activity in the right-hand compartment is 1.00, the cell potential is the standard electrode potential of the Ag+/Ag half-reaction.

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This galvanic cell develops a potential of 0.799 V with the silver electrode

the standard electrode potential is given a positive sign, and we write

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18C-5 Effect of Concentration on Electrode Potentials: The Nernst Equation

Consider the reversible half-reaction

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E0 = the standard electrode potential, which is characteristic for each half-reaction

R = the ideal gas constant, 8.314 J K － 1 mol － 1

T = temperature, K

n = number of moles of electrons that appears in the half-reaction for the electrode process as writtenF = the faraday 96,485 C (coulombs) per mole of electronsIf we substitute numerical values for the constants, convert to base 10 logarithms, and specify 25°C for the temperature, we get 509

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Nernst equation

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If we substitute numerical values for the constants, convert to base 10 logarithms, and specify 25°C for the temperature, we get

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1. The standard electrode potential is a relative quantity in the sense that it is the potential of an electrochemical cell in which the reference electrode is the standard hydrogen electrode, whose potential has been assigned a value of 0.000 V.

2. The standard electrode potential for a half-reaction refers exclusively to a reduction reaction;

3. The standard electrode potential measures the relative force tending to drive the half-reaction from the reactants and products are at their equilibrium activities

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18C-6 The Standard Electrode Potential, E0

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4. The standard electrode potential is independent of the number of moles of reactant and product shown in the balanced half-reaction.

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5. A positive electrode potential indicates that the half-reaction in question is spontaneous with respect to the standard hydrogen electrode half-reaction.

6. The standard electrode potential for a half-reaction is temperature dependent.

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System involving precipitates or complex ions

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Ch 19 Applications of Standard Electrode Potentials

EXAMPLE 19-1 Calculate the thermodynamic potential of the following

cell and the free energy change associated with the cell reaction.

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2e- + Fe2+ Fe

2Ag 2Ag+ + 2e-Oxidation:

Reduction:

What is the equilibrium constant for the following reaction at 250C? Fe2+ (aq) + 2Ag (s) Fe (s) + 2Ag+ (aq)

=0.0257 V

nln KEcell

0

19.4

n = 2

EXAMPLE 19-2 Calculate the potential of the cell

Ag Ag+ ( 0.0200 M) (0.0200M) Cu2+ Cu

EXAMPLE 19-3 Calculate the potential of the following cell and indicate

the reaction that would occur spontaneously if the cell were short circuited (Figure 19-1).

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EXAMPLE 19-4 Calculate the cell potential for

Note that this cell does not require two compartments (nor a salt bridge) because molecular H2 has little tendency to react directly with the low concentration of Ag ＋ in the electrolyte solution. This is an example of a cell without liquid junction (Figure 19-2).

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EXAMPLE 19-5

Calculate the potential for the following cell using (a) concentration (b) activity

where x = 5.00x10-4, 2.00x10-3, 1.00x10-2, and 5.00x10-2 Zn ZnSO4 ( xM), PbSO4 (sat'd) Pb

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PbSO4(s) + 2e Pb(s) + SO42- E0 = - 0.350 VPbSO4/Pb

(a) concentration

EXAMPLE 19-5

Calculate the potential for the following cell using (a) concentration (b) activity

where x = 5.00x10-4, 2.00x10-3, 1.00x10-2, and 5.00x10-2 Zn ZnSO4 ( xM), PbSO4 (sat'd) Pb

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(b) activity 活性

EXAMPLE 19-6

Calculate the potential required to initiate deposition of copper from a solution that is 0.010 M in CuSO

4 and contains sufficient H2SO4 to give a pH of 4.00.

The deposition of copper necessarily occurs at the cathode.

Since there is no more easily oxidizable species than water in the system, O2 will evolve at the anode.

EXAMPLE 19-7 D. A. MacInnes found that a cell similar to that sho

wn in Figure 19-2 had a potential of 0.52053 V. The cell is described by the following notation.

Calculate the standard electrode potential for the half-reaction (by activities)

19C CALCULATING REDOX EQUILIBRIUM CONSTANTS （氧化還原反應的平衡常數） Thus, at chemical equilibrium, we may write

or

We can generalize Equation 19-6 by stating that at equilibrium, the electrode potentials for all half-reactions in an oxidation/reduction system are equal.

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Cu(s) + 2Ag+ 2Ag(s) + Cu2+

EXAMPLE 19-8

Calculate the equilibrium constant for the reaction shown in Equation 19-4 at 25°C.

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EXAMPLE 19-9 Calculate the equilibrium constant for the reaction

2Fe3+ + 3I- 2Fe2+ + I3-

2 Fe3+ + 2e 2 Fe2+ E0 = 0.771V

I3- + 2e 3 I- E0 = 0.536V

EXAMPLE 19-10 Calculate the equilibrium constant for the

reaction

Again we have multiplied both equations by integers so that the numbers of electrons are equal. When this system is at equilibrium.

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