Transcript Document 7549853

CHAPTER
Introduction to Electrochemistry
Chapter18 p
 Oxidation ：氧化反應
 Reduction：還原反應
 Reducing agent : 還原劑
還原其他物種，自身進行氧化反應
 Oxidizing agent ：氧化劑
氧化其他物種，自身進行還原反應
Chapter18 p
Example 18-1
The following reactions are spontaneous and thus
proceed to the right, as written
What can we deduce regarding the strengths of H+,
Ag+, Cd2+, Zn2+ as electron acceptors? (or oxidizing
agents)
2H+ + Cd(s)
H2 + Cd2+
2Ag+ + H2(g)
2Ag(s) + 2H+
Cd2+ + Zn(s)
Cd(s) + Zn2+
Ag+ > H+ > Cd2+ > Zn2+
493
銅片浸入硝酸銀水溶液中
Ag+
+ e
Ag(s)
Figure 18-1
Photograph of a “ silver tree.”
Chapter18 p 493
Galvanic Cell 賈法尼電池
anode
oxidation
cathode
reduction
spontaneous
redox reaction
Chapter18 p
19.2
Figure 18-2
(a)A galvanic cell at open circuit;
Chapter18 p 495
(b) a galvanic cell doing work;
Chapter18 p 495
(c) an electronlytic cell. 電解電池
Chapter18 p 495
18B-3 Representing Cells Schematically
電池表示方法
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 liquidjunction potential develops at each of these interfaces.
Chapter18 p 498
陰極(cathode)
陽極(anode)
Chapter18 p
Figure 18-3
Movement of
charge in a
galvanic cell.
Chapter18 p 500
18C Electrode Potentials
電極
電位
The cell potential Ecell is related to the free energy of
the reaction ΔG by
Chapter18 p 499
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.
Chapter18 p 500
(a)
Chapter18 p 501
(b)
Chapter18 p 501
(c)
Chapter18 p 501
Figure 18-5
Cell 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.
Chapter18 p 503
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
Chapter18 p
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
Chapter18 p 504
Figure 18-6
The hydrogen gas electrode.
By convention, the
potential of the
standard hydrogen
electrode is assigned a
value of 0.000 V at all
temperatures.
Chapter18 p 505
Standard Electrode Potentials 標準電極電位
Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)
Anode (oxidation):
Zn (s)
Zn2+ (1 M) + 2e-
Cathode (reduction): 2e- + 2H+ (1 M)
Zn (s) + 2H+ (1 M)
Zn2+ + H2 (1 atm)
H2 (1 atm)
Chapter18 p
19.3
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.
Chapter18 p
p.505
The standard electrode potential, E0, of a halfreaction 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
Chapter18 p 506
Figure 18-7
Measurement of the
electrode potential
for an Ag electrode.
If the silver ion
activity in the righthand compartment is
1.00, the cell
potential is the
standard electrode
potential of the
Ag+/Ag halfreaction.
Chapter18 p 506
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
Chapter18 p 506
18C-5 Effect of Concentration on Electrode Potentials:
The Nernst Equation
Consider the reversible half-reaction
Chapter18 p 508
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 written
F = the faraday 96,485 C (coulombs) per mole of
electrons
If we substitute numerical values for the constants,
convert to base 10 logarithms, and specify 25°C for the
temperature, we get
Chapter18 p 509
If we substitute numerical values for the constants,
convert to base 10 logarithms, and specify 25°C for the
temperature, we get
Nernst equation
Chapter18 p
p.509
18C-6 The Standard Electrode Potential, E0
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
4. The standard electrode potential is independent of the number
of moles of reactant and product shown in the balanced halfreaction.
Chapter18 p 511
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5. A positive electrode potential indicates that the halfreaction 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.
Chapter18 p
Chapter18 p 512
System involving precipitates or complex ions
Chapter18 p
Ch 19 Applications of Standard
Electrode Potentials
Chapter18 p
EXAMPLE 19-1 計算下列電池的電位與自由能變化量
 Calculate the thermodynamic potential of the following
cell and the free energy change associated with the cell
reaction.
395
What is the equilibrium constant for the following reaction
at 250C? Fe2+ (aq) + 2Ag (s)
Fe (s) + 2Ag+ (aq)
0
Ecell
=
0.0257 V
ln K
n
Oxidation:
Reduction:
2e-
+
2Ag
2Ag+ + 2e-
Fe2+
Fe
n=2
19.4
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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計算下列電池的電位
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 計算下列電池的電位
Calculate the potential for the following cell using (a)
concentration (b) activity
Zn ZnSO4 ( xM), PbSO4 (sat'd) Pb
where x = 5.00x10-4, 2.00x10-3, 1.00x10-2, and 5.00x10-2
(a) concentration
PbSO4(s) + 2e
Pb(s) + SO42-
E0PbSO /Pb = - 0.350 V
4
計算下列電池的電位
EXAMPLE 19-5
Calculate the potential for the following cell using (a)
concentration (b) activity
Zn ZnSO4 ( xM), PbSO4 (sat'd) Pb
where x = 5.00x10-4, 2.00x10-3, 1.00x10-2, and 5.00x10-2
(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 CuSO4 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
shown 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)
Cu(s) + 2Ag+
2Ag(s) + Cu2+
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 halfreactions in an oxidation/reduction system are equal.
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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
3+
2Fe
-
+ 3I
2 Fe3+ + 2e
I3
-
+
2+
2e
2Fe
+
I3
2 Fe2+ E0 = 0.771V
-
3I
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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