Introduction to Electrochemistry

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CHAPTER
18
Introduction to Electrochemistry
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• Oxidation ????
• Reduction????
• Reducing agent ???
• ??????,????????
• Oxidizing agent ???
• ??????,????????

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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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???????????
Figure 18-1 Photograph of a silver tree.
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Galvanic Cell ?????
anode oxidation
cathode reduction
spontaneous redox reaction
19.2
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Figure 18-2 (a)A galvanic cell at open circuit
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(b) a galvanic cell doing work
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(c) an electronlytic cell. ????
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??????
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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??(cathode)
??(anode)
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Figure 18-3 Movement of charge in a galvanic
cell.
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18C Electrode Potentials ??
??

• The cell potential Ecell is related to the free
  energy of the reaction ?G 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 18-5 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.
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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 18-6 The hydrogen gas electrode.
By convention, the potential of the standard
hydrogen electrode is assigned a value of 0.000 V
at all temperatures.
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Standard Electrode Potentials ??????
Zn (s) Zn2 (1 M) H (1 M) H2 (1 atm) Pt
(s)
Anode (oxidation)
Cathode (reduction)
19.3
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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 18-7 Measurement 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
  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 25C for the temperature, we get

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

Nernst equation
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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.

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

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

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

1. 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
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????????????????

• 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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Oxidation
n 2
Reduction
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• EXAMPLE 19-2
• Calculate the potential of the cell

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• 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
?????????
(a) concentration
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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
(b) activity ??
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• 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.

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• 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)

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19C CALCULATING REDOX EQUILIBRIUM
CONSTANTS(???????????)

• 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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EXAMPLE 19-8 Calculate the equilibrium
constant for the reaction shown in Equation 19-4
at 25C.
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EXAMPLE 19-9 Calculate the equilibrium
constant for the reaction
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• 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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