1 1 Electrochemistry
Welcome message from author
This document is posted to help you gain knowledge. Please leave a comment to let me know what you think about it! Share it to your friends and learn new things together.
Transcript
1
1
Electrochemistry
2
2
Half-Reactions
1. Balancing Oxidation–Reduction Reactions in Acidic and Basic Solutions
Voltaic Cells
2. Construction of Voltaic Cells
3. Notation for Voltaic Cells
4. Cell Potential
5. Standard Cell Potentials and Standard Electrode Potentials
3
3
6. Equilibrium Constants from Cell Potentials
7. Dependence of Cell Potentials on Concentration
8. Some Commercial Voltaic Cells
Electrolytic Cells
9. Electrolysis of Molten Salts
10. Aqueous Electrolysis
11. Stoichiometry of Electrolysis
4
4
Electrochemistry
• Balancing Oxidation–Reduction Reactions in Acidic and Basic Solutions
a. Learn the steps for balancing oxidation–reduction reactions using the half-reaction method.
b. Balance equations by the half-reaction method (acidic solutions).
c. Learn the additional steps for balancing oxidation–reduction reactions in basic solution using the half-reaction method.
d. Balance equations using the half-reaction method (basic solution).
Learning Objectives
5
5
2. Construction of Voltaic Cells
– Define electrochemical cell, voltaic (galvanic
cell), electrolytic cell, and half-cell.
– Describe the function of the salt bridge in a
voltaic cell.
– State the reaction that occurs at the anode and
the cathode in an electrochemical cell.
– Define cell reaction.
– Sketch and label a voltaic cell.
6
6
3. Notation for Voltaic Cells
– Write the cell reaction from the cell notation.
4. Cell Potential
– Define cell potential and volt.
– Calculate the quantity of work from a given
amount of cell reactant.
7
7
5. Standard Cell Potentials and Standard Electrode Potentials
– Explain how the electrode potential of a cell is an intensive property.
– Define standard cell potential and standard electrode potential.
– Interpret the table of standard reduction potentials.
– Determine the relative strengths of oxidizing and reducing agents.
– Determine the direction of spontaneity from electrode potentials.
– Calculate cell potentials from standard potentials.
8
8
6. Equilibrium Constants from Cell Potentials
– Calculate the free-energy change from electrode
potentials.
– Calculate the cell potential from free-energy
change.
– Calculate the equilibrium constant from cell
potential.
7. Dependence of Cell Potential on Concentration
– Calculate the cell potential for nonstandard
conditions.
– Describe how pH can be determined using a
glass electrode.
9
9
8. Some Commercial Voltaic Cells
– Describe the construction and reactions of a
zinc–carbon dry cell, a lithium–iodine battery, a
lead storage cell, and a nickel–cadmium cell.
– Explain the operation of a proton-exchange
membrane fuel cell.
– Explain the electrochemical process of the
rusting of iron.
– Define cathodic protection.
9. Electrolysis of Molten Salts
– Define electrolysis.
10
10
10. Aqueous Electrolysis
– Learn the half-reactions for water undergoing
oxidation and reduction.
– Predict the half-reactions in an aqueous
electrolysis.
11. Stoichiometry of Electrolysis
– Calculate the amount of charge from the amount
of product in an electrolysis.
– Calculate the amount of product from the
amount of charge in an electrolysis.
11
11
Our first step in studying electrochemical cell is to
balance its oxidation–reduction reaction.
We will use the half-reaction method from chapter 5 and
extend it to acidic or basic solutions.
In this chapter we will focus on electron transfer rather
than proton transfer so the hydronium ion, H3O+(aq), will
be represented by its simpler notation, H+(aq). Only the
notation, not the chemistry, is different.
12
12
Oxidation-Reduction Reactions
• In Chapter 5 we introduced the half-reaction
method for balancing simple oxidation-reduction
reactions.
– Oxidation-reduction reactions always involve a transfer of electrons from one
species to another.
– Recall that the species losing electrons is oxidized, while the species gaining
electrons is reduced.
13
13
Oxidation-Reduction Reactions
• Describing Oxidation-Reduction Reactions
– An oxidizing agent is a species that oxidizes another species; it is itself reduced.
– A reducing agent is a species that reduces
another species; it is itself oxidized.
)()(2
)(2
)( saqaqs CuFeCuFe ++++→→→→++++++++++++
oxidizing agent
reducing agentLoss of 2 e-1 oxidation
Gain of 2 e-1 reduction
14
14
What reaction is going
on and what is the
equation?
A voltaic cell employs a spontaneous oxidation–reduction reaction
as a source of energy. It separates the reaction into two half-
reactions, physically separating one half-reaction from the other.
15
15
We will be studying more complex situations, so our initial
analysis is key.
First, we need to identify what is being oxidized and what
is being reduced.
Then, we determine if the reaction is in acidic or basic
conditions.
16
16
Balancing Oxidation-Reduction Equations in Acid Solution
1. Assign oxidation numbers
2. Determine half reactions
3. Complete and balance each half-reaction
a. Balance all atoms except O and H
b. Balance O atoms by adding H2O’s to one side
c. Balance H atoms by adding H+ to one side
d. Balance electric charges by adding e- to the more positive side
4. Combine each half reaction to obtain the final balanced
oxidation reduction equation (Multiply by appropriate
factors and cancel species appearing on both sides of equation
17
17
Balancing Oxidation-Reduction Equations in Basic Solution
1. Assign oxidation numbers
2. Determine half reactions
3. Complete and balance each half-reaction
a. Balance all atoms except O and H
b. Balance O atoms by adding H2O’s to one side
c. Balance H atoms by adding H+ to one side
d. Balance electric charges by adding e- to the more positive side
4. Note the number of H+ ions in the equation. Add this number
of OH- ions to both sides of the equation.
6. Combine each half reaction to obtain the final balanced
oxidation reduction equation (Multiply by appropriate
factors and cancel species appearing on both sides of equation
5. Simplify the equation by noting that H+ and OH- react to form
H2O. Cancel waters that occur on both sides of the equation and
reduce the equation to simplest terms.
18
18
Balance in acid solution:
Mn+2 + BiO3- → MnO4
- + Bi+3
Balance in Basic solution:
S2- + I2 → SO42- + I-
19
19
Balance in acid solution:
Mn+2 + BiO3- → MnO4
- + Bi+3
2(Mn+2 + 4 H2O → MnO4- + 8 H+ + 5e- )
5(6 H+ + BiO3- + 2e- → Bi+3 + 3 H2O)
2Mn+2 + 8 H2O → 2MnO4- + 16 H+ + 10e- )
30 H+ + 5 BiO3- + 10e- → 5Bi+3 + 15 H2O)
2Mn2+ + 14 H+ + 5 BiO3- → 2 MnO4
- + 5 Bi3+ + 7 H2O
20
20
Balance in Basic solution:
S2- + I2 → SO42- + I-
8 OH- +4H2O + S2- → SO42- + 8 H+ + 8 OH-+ 8e-
I2 + 2 e- → 2I-
4 (I2 + 2 e- → 2I-)
4 I2 + 8 e- → 8I-
8 OH- + S2- + 4 I2 → SO42- + 8 I- +4 H2O
8 OH- +4H2O + S2- → SO42- + 8 H+ + 8 OH-+ 8e-
8 OH- +4H2O + S2- → SO42- + 8 H2O + 8e-
21
21
• Dichromate ion in acidic solution is an oxidizing
agent. When it reacts with zinc, the metal is oxidized
to Zn2+, and nitrate is reduced. Assume that
dichromate ion is reduced to Cr3+.
• Write the balanced ionic equation for this reaction
using the half-reaction method.
22
22
First we determine the oxidation numbers of N and Zn:
+6 +3 0 +2
Cr2O72-(aq) � Cr3+(aq) and Zn(s) � Zn2+(aq)
Cr was reduced from +6 to +3.
Zn was oxidized from 0 to +2.
Now we balance the half-reactions.
23
23
Oxidation half-reaction
Zn(s) � Zn2+(aq) + 2e-
Reduction half-reaction
First we balance Cr and O:
Cr2O72-(aq) � 2Cr3+(aq) + 7H2O(l)
Next we balance H:
14H+(aq) + Cr2O72-(aq) � 2Cr3+(aq) + 7H2O(l)
Then we balance e-:
6e- + 14H+(aq) + Cr2O72-(aq) � 2Cr3+(aq) + 7H2O(l)
24
24
Now we combine the two half-reactions by multiplying
the oxidation half reaction by 3.
3Zn(s) � 3Zn2+(aq) + 6e-
6e- + 14H+(aq) + Cr2O72-(aq) � 2Cr3+(aq) + 7H2O(l)
3Zn(s) + 14H+(aq) + Cr2O72-(aq) �
3Zn2+(aq) + 2Cr3+(aq) + 7H2O(l)
Check that atoms and charge are balanced.
3Zn; 14H; 2Cr; 7O; charge +12
25
25
To balance a reaction in basic conditions, first follow
the same procedure as for acidic solution.
• Then
1. Add one OH- to both sides for each H+.
2. When H+ and OH- occur on the same side,
combine them to form H2O.
3. Cancel water molecules that occur on both sides.
Balance in Basic solution:
S2- + I2 → SO42- + I-
Did This Earlier
26
26
• Lead(II) ion, Pb2+, yields the plumbite ion, Pb(OH)3-,
in basic solution. In turn, this ion is oxidized in basic
hypochlorite solution, ClO-, to lead(IV) oxide, PbO2.
• Balance the equation for this reaction using the half-
reaction method. The skeleton equation is
• Pb(OH)3- + ClO- � PbO2 + Cl-
Another example
27
27
First we determine the oxidation number of Pb and
Cl in each species.
+2 +1 +4 -1
Pb(OH)3- + ClO- � PbO2 + Cl-
Pb is oxidized from +2 to +4.
Cl is reduced from +1 to -1.
Now we balance the half-reactions.
28
28
Oxidation half-reaction
First we balance Pb and O:
Pb(OH)3-(aq) � PbO2(s) + H2O(l)
Next we balance H:
Pb(OH)3-(aq) � PbO2(s) + H2O(l) + H+(aq)
Then we balance e-:
Pb(OH)3-(aq) � PbO2(s) + H2O(l) + H+(aq) + 2e-
29
29
Reduction half-reaction
First we balance Cl and O:
ClO-(aq) � Cl-(aq) + H2O(l)
Next we balance H:
2H+(aq) + ClO-(aq) � Cl-(aq) + H2O(l)
Finally we balance e-:
2e- + 2H+(aq) + ClO-(aq) � Cl-(aq) + H2O(l)
30
30
Now we combine the half-reactions.
Pb(OH)3-(aq) � PbO2(s) + H2O(l) + H+(aq) + 2e-
2e- + 2H+(aq) + ClO-(aq) � Cl-(aq) + H2O(l)
H+(aq) + Pb(OH)3-(aq) + ClO-(aq) �
PbO2(s) + Cl-(aq) + 2H2O(l)
31
31
H+(aq) + Pb(OH)3-(aq) + ClO-(aq) �
PbO2(s) + Cl-(aq) + 2H2O(l)
To convert to basic solution, we add OH- to each side,
converting H+ to H2O.
H2O(l) + Pb(OH)3-(aq) + ClO-(aq) �
PbO2(s) + Cl-(aq) + 2H2O(l) + OH-(aq)
Finally, we cancel the H2O that is on both sides.
Pb(OH)3-(aq) + ClO-(aq) �
PbO2(s) + Cl-(aq) + H2O(l) + OH-(aq)
Do exercise 19.1 & 2 and see problems 19.35-37
32
32
The next several topics describe battery cells or
voltaic cells (galvanic cells).
An electrochemical cell is a system consisting of
electrodes that dip into an electrolyte and in which a
chemical reaction either uses or generates an electric
current.
A voltaic or galvanic cell is an electrochemical cell
in which a spontaneous reaction generates an electric
current.
An electrolytic cell is an electrochemical cell in
which an electric current drives an otherwise
nonspontaneous reaction.
33
33
Oxidation-Reduction Reactions
• In this chapter we will show how a cell is
constructed to physically separate an oxidation-reduction reaction into two half-reactions.
– The force with which electrons travel from
the oxidation half-reaction to the reduction
half-reaction is measured as voltage.
34
34
Electrochemistry• An electrochemical cell is a system
consisting of electrodes that dip into an electrolyte in which a chemical reaction either uses or generates an electric current.
– A voltaic, or galvanic, cell is an electrochemical cell in which a spontaneous reaction generates an electric current.
– An electrolytic cell is an electrochemical cell in which an electric current drives an otherwise nonspontaneous reaction.
– In this chapter we will discuss the basic principles behind these cells and explore some of their commercial uses.
35
35
Voltaic Cells
• A voltaic cell consists of two half-cells that are
electrically connected.
– Each half-cell is a portion of the electrochemical cell in which a half-reaction takes place.
– A simple half-cell can be made from a metal strip dipped into a solution of its metal ion.
– For example, the zinc-zinc ion half cell consists consists of a zinc strip dipped into a solution of a zinc salt.
36
36
Voltaic Cells
• A voltaic cell consists of two half-cells
that are electrically connected.
– Another simple half-cell consists of a copper strip dipped into a solution of a copper salt.
– In a voltaic cell, two half-cells are connected in such a way that electrons flow from one metal electrode to the other through an external circuit.
– Figure 19.2 illustrates an atomic view of a zinc/copper voltaic cell.
37
37
Figure 19.2: Atomic view of voltaic cell.
38
38
x
The zinc metal atom loses two electrons, forming Zn2+
ions.
The Cu2+ ions gain two electrons, forming solid
copper.
The electrons flow through the external circuit from the zinc electrode to the copper electrode.
Ions flow through the salt bridge to maintain charge balance.
39
39
Voltaic Cells
• As long as there is an external circuit, electrons
can flow through it from one electrode to the
other.
– Because zinc has a greater tendency to lose electrons than copper, zinc atoms in the zinc electrode lose electrons to form zinc ions.
– The electrons flow through the external circuit to the copper electrode where copper ions gain the electrons to become copper metal.
40
40
Voltaic Cells
• The two half-cells must also be connected
internally to allow ions to flow between them.
– Without this internal connection, too much positive charge builds up in the zinc half-cell (and too much negative charge in the copper half-cell) causing the reaction to stop.
– Figure 19.3A and 19.3B show the two half-cells of a voltaic cell connected by salt bridge.
41
41
Figure 19.3: Two electrodes are connected by an
external circuit.
42
42
Voltaic Cells
• A salt bridge is a tube of an electrolyte in a gel
that is connected to the two half-cells of a voltaic
cell.
– The salt bridge allows the flow of ions but
prevents the mixing of the different solutions that would allow direct reaction of the cell reactants.
– Figure 19.3C shows an actual setup of the
zinc-copper cell.
43
43
44
44
Voltaic Cells• The two half-cell reactions, as noted earlier, are:
– The first reaction, in which electrons are lost, is the oxidation half-reaction.
– The electrode at which oxidation occurs is
the anode.
−−−−++++ ++++→→→→ e2)aq(Zn)s(Zn 2
)s(Cue2)aq(Cu2
→→→→++++−−−−++++
(oxidation half-reaction)
(reduction half-reaction)
– The second reaction, in which electrons
are gained, is the reduction half-reaction.
– The electrode at which reduction occurs is the cathode.
45
45
Voltaic Cells
• Note that the sum of the two half-reactions
– Note that electrons are given up at the anode and thus flow from it to the cathode where reduction occurs.
)s(Cu)aq(Zn)aq(Cu)s(Zn22
++++→→→→++++++++++++
is the net reaction that occurs in the
voltaic cell; it is called the cell reaction
– The anode in a voltaic cell has a negative sign because electrons flow from it. (See Figure 20.4 andAnimation: Anode Reaction)
– The cathode in a voltaic cell has a positive sign (See Animation: Cathode Reaction)
46
46
Figure 19.4:
Voltaic Cell
Do exercise 19.3
And look at
Problems 19.43-44
47
47
Notation for Voltaic Cells
• It is convenient to have a shorthand way of
designating particular voltaic cells.
– The anode (oxidation half-cell) is written on the left. The cathode (reduction half-cell) is written on the right.
– The cell consisting of the zinc-zinc ion half-cell and the copper-copper ion half-cell, is written
)s(Cu|)aq(Cu||)aq(Zn|)s(Zn22 ++++++++
anode cathode
48
48
– The cell terminals are at the extreme ends in the cell notation.
)s(Cu|)aq(Cu||)aq(Zn|)s(Zn22 ++++++++
anode cathodesalt bridge
– The cell consisting of the zinc-zinc ion half-cell and the copper-copper ion half-cell, is written
– A single vertical bar indicates a phase boundary, such as between a solid terminal and the electrode solution.
anode cathodesalt bridge
)s(Cu|)aq(Cu||)aq(Zn|)s(Zn22 ++++++++
49
49
Notation for Voltaic Cells
• When the half-reaction involves a gas, an
inert material such as platinum serves as a
terminal and an electrode surface on which
the reaction occurs.
– Figure 19.5 shows a hydrogen electrode; hydrogen bubbles over a platinum plate immersed in an acidic solution.
– The cathode half-reaction is
)g(He2)aq(H2 2→→→→++++
−−−−++++
50
50
– The notation for the hydrogen electrode, written as a cathode, is
Pt|)g(H|)aq(H 2
++++
– To write such an electrode as an anode,
you simply reverse the notation.
)aq(H|)g(H|Pt 2
++++
– In the cell notation, these are written in parentheses. For example,
Pt|)atm0.1(H|)aq(H||)M0.1(Zn|)s(Zn 22 ++++++++
51
51
Figure 19.5:
Hydrogen
Electrode
52
52
A Problem To Consider
• Give the overall cell reaction for the voltaic cell
Pt|)atm0.1(H|)aq(H||)M0.1(Cd|)s(Cd 22 ++++++++
– The half-cell reactions are
)g(He2)aq(H2 2→→→→++++
−−−−++++
−−−−++++++++→→→→ e2)aq(Cd)s(Cd
2
)g(H)aq(Cd)aq(H2)s(Cd 22
++++→→→→++++++++++++
Do Exercise 19.4 and see problems 19.49-50
53
53
Electromotive Force (Cell Potential)
• The movement of electrons is analogous to the
pumping of water from one point to another.
– Water moves from a point of high pressure to a point of lower pressure. Thus, a pressure difference is required.
– The work expended in moving the water through a pipe depends on the volume of water and the pressure difference.
– An electric charge moves from a point of high electrical potential (high electrical pressure) to one of lower electrical potential.
– The work expended in moving the electrical charge through a conductor depends on the amount of charge and the potential difference.
54
54
Electromotive Force• Potential difference is the difference in electric
potential (electrical pressure) between two points.
– You measure this quantity in volts.
– The volt, V, is the SI unit of potential difference equivalent to 1 joule of energy
per coulomb of charge.
CJ1volt1 ====
Electrical work = charge x potential difference
Joules = coulombs x volts
55
55
Electromotive Force
• The Faraday constant, F, is the magnitude of
charge on one mole of electrons; it equals 96,500
coulombs (9.65 x 104 C).
– In moving 1 mol of electrons through a circuit, the numerical value of the work done by a voltaic cellis the product of the Faraday constant (F) times the potential difference between the electrodes.
)coulombvolts(J/)F(coulombswork(J) ××××−−−−====
work done by the system
56
56
– In the normal operation of a voltaic cell, the potential difference (voltage) across the electrodes is less than than the maximum possible voltage of the cell.
Electromotive Force
– The actual flow of electrons reduces the electrical pressure.
– Thus, a cell voltage has its maximum value when no current flows.
57
57
Electromotive Force
• The maximum potential difference between
the electrodes of a voltaic cell is referred to
as the electromotive force (emf) of the cell,
or Ecell.
– It can be measured by an electronic digital voltmeter (Figure 19.6), which draws
negligible current.
58
58
59
59
Electromotive Force
• We can now write an expression for the
maximum work attainable by a voltaic
cell.
– The maximum work for molar amounts of
reactants is
cellmax nFE−−−−====w
– Let n be the number of (mol) electrons transferred in the overall cell reaction.
Work Exercise 19.6 Look at Problems 19.57-60
60
60
A Problem To Consider• The emf of the electrochemical cell below is 0.650
V. Calculate the maximum electrical work of this
cell when 0.500 g H2 is consumed.
)aq(H2)l(Hg2)g(H)aq(Hg 2
2
2
++++++++++++++++
– The half-reactions are
)l(Hg2e2)aq(Hg2
2
−−−−++++++++
−−−−++++++++ e2)aq(H2)g(H2
– n = 2, and the maximum work for the reaction is written as
cellmax nFE−−−−====w
)V650.0()C1065.9(2 4max
××××××××××××−−−−====w
61
61
)V650.0()C1065.9(2 4max
××××××××××××−−−−====w
cellmax nFE−−−−====w
J1025.1 5max
××××−−−−====w
• The emf of the electrochemical cell below is 0.650
V. Calculate the maximum electrical work of this
cell when 0.500 g H2 is consumed.
– For 0.500 g H2, the maximum work is
J1009.3Hmol1
J1025.1
Hg02.2
Hmol1Hg500.0
4
2
5
2
22
××××−−−−====××××−−−−
××××××××
)aq(H2)l(Hg2)g(H)aq(Hg 2
2
2
++++++++++++++++
62
62
Standard Cell emf’s and Standard Electrode
Potentials
• A cell emf is a measure of the driving force
of the cell reaction.
– The reaction at the anode has a definite oxidation potential, while the reaction at the cathode has a definite reduction potential.
– Thus, the overall cell emf is a combination of these two potentials.
–Ecell = oxidation potential + reduction potential
63
63
– A reduction potential is a measure of the tendency to gain electrons in the reduction half-reaction.
– You can look at the oxidation half-reaction
as the reverse of a corresponding reduction reaction.
– The oxidation potential for an oxidation half-reaction is the negative of the
reduction potential for the reverse reaction.
64
64
– Consider the zinc-copper cell described earlier.
)s(Cu|)aq(Cu||)aq(Zn|)s(Zn22 ++++++++
Standard Cell emf’s and Standard Electrode
Potentials
• By convention, the Table of Standard
Electrode Potentials (Table 19.1) are
tabulated as reduction potentials.
−−−−++++++++→→→→ e2)aq(Zn)s(Zn
2
– The two half-reactions are
)s(Cue2)aq(Cu2
→→→→++++−−−−++++
65
65
– The zinc half-reaction is an oxidation.
– If you write EZn for the reduction potential of zinc,
then –EZn is the oxidation potential of zinc.
)s(Zne2)aq(Zn2
→→→→++++−−−−++++
−−−−++++++++→→→→ e2)aq(Zn)s(Zn
2
(EZn) Reduction
-(EZn)
– The copper half-reaction is a reduction..
(ECu))s(Cue2)aq(Cu2
→→→→++++−−−−++++
– Write ECu for the electrode potential.
– For this cell, the cell emf is the sum of the reduction potential for the copper half-cell and the oxidation potential for the zinc half-cell.
Oxidat.
66
66
)E(EE ZnCucell−−−−++++====
ZnCucell EEE −−−−====
– Note that the cell emf is the difference between
the two electrode potentials.
anodecathodecellEEE −−−−====
– In general, Ecell is obtained by subtracting the anode potential from the cathode potential.
– The electrode potential is an intensive property whose value is independent of the amount of species in the reaction.
67
67
– Thus, the electrode potential for the half-reaction
)s(Cu2e4)aq(Cu22
→→→→++++−−−−++++
is the same as for
)s(Cue2)aq(Cu2
→→→→++++−−−−++++
68
68
– Note that individual electrode potentials require
that we choose a reference electrode.
Tabulating Standard Electrode Potentials
• The standard emf, E°cell, is the emf of a cell
operating under standard conditions of
concentration (1 M), pressure (1 atm), and
temperature (25°C).
– You arbitrarily assign this reference electrode a
potential of zero and obtain the potentials of the
other electrodes by measuring the emf’s.
69
69
– By convention, the reference chosen for comparing electrode potentials is the standard
hydrogen electrode. (see Figure 19.5)
– Standard electrode potentials (Table 19.1) are
measured relative to this hydrogen reference.
70
70
– For example, when you measure the emf of a
cell composed of a zinc electrode connected to a hydrogen electrode, you obtain 0.76 V.
Tabulating Standard Electrode Potentials
• The standard electrode potential, E°, is the
electrode potential when concentrations of solutes
are 1 M, gas pressures are 1 atm, and the
temperature is 25°C. (Table 19.1)
– Since zinc acts as the anode (oxidation) in this
cell, its reduction potential is listed as –0.76 V.
71
71
72
72
73
73
74
74
Strengths of Oxidizing and Reducing Agents
• Standard electrode potentials are useful in determining the
strengths of oxidizing and reducing agents under
standard-state conditions.
– A reduction half-reaction has the general form
speciesreducednespeciesoxidized →→→→++++−−−−
– The oxidized species acts as an oxidizing agent.
– Consequently, the strongest oxidizing agents in a table of standard electrode potentials are the oxidized speciescorresponding to the half-reactions with the largest (most positive) Eo values. (For example F2(g))
75
75
– An oxidation half-reaction has the general form
−−−−++++→→→→ nespeciesoxidizedspeciesreduced
– The reduced species acts as a reducing agent.
– Consequently, the strongest reducing agentsin a table of standard electrode potentials are the reduced species corresponding to the half-reactions with the smallest (most negative) Eo values. (for example, Li(s))
76
76
• Which is the stronger reducing agent under
standard conditions: Sn2+ (to Sn4+) or Fe (to
Fe2+)?
• Which is the stronger oxidizing agent under
standard conditions: Cl2 or MnO4-?
The stronger reducing agent will be oxidized and has
the more negative electrode potential.
The standard (reduction) potentials areSn2+ to Sn4+ E = 0.15 VFe to Fe2+ E = –0.41 V
The stronger reducing agent is Fe (to Fe2+).
77
77
• The stronger oxidizing agent will be reduced.
• The standard (reduction) potentials are
• Cl2 to Cl- E = 1.36 V
• MnO4- to Mn2+ E = 1.49 V
• The stronger oxidizing agent is MnO4-.
78
78
Calculating Cell emf’s from Standard Potentials
• The emf of a voltaic cell constructed from standard
electrodes is easily calculated using a table of electrode
potentials.– Consider a cell constructed of the following two half-
reactions.
V0.40E);s(Cde2)aq(Cdo2
−−−−====→→→→++++−−−−++++
V0.80E);s(Age1)aq(Ago
====→→→→++++−−−−++++
– You will need to reverse one of these reactions to obtain the oxidation part of the cell reaction.
V0.40E);s(Cde2)aq(Cdo2
−−−−====→→→→++++−−−−++++
V0.80E);s(Age1)aq(Ago
====→→→→++++−−−−++++
79
79
– This will be Cd, because has the more negative electrode potential.
V0.40E);s(Cde2)aq(Cdo2
−−−−====→→→→++++−−−−++++
V0.80E);s(Age1)aq(Ago
====→→→→++++−−−−++++
– Therefore, you reverse the half-reaction and change the sign of the half-cell potential.
V0.40E;e2)aq(Cd)s(Cdo2
====++++→→→→−−−−++++
V0.80E);s(Age1)aq(Ago
====→→→→++++−−−−++++
– We must double the silver half-reaction so that when the reactions are added, the electrons
cancel.
80
80
V0.40E;e2)aq(Cd)s(Cdo2
====++++→→→→−−−−++++
V0.80E);s(Age1)aq(Ago
====→→→→++++−−−−++++
– This does not affect the half-cell potentials, which do not depend on the amount of substance.
V0.40E;e2)aq(Cd)s(Cdo2
====++++→→→→−−−−++++
V0.80E);s(2Age2)aq(2Ago
====→→→→++++−−−−++++
– Now we can add the two half-reactions to obtain the overall cell reaction and cell emf.
V0.40E;e2)aq(Cd)s(Cdo2
====++++→→→→−−−−++++
V0.80E);s(Ag2e2)aq(Ag2o
====→→→→++++−−−−++++
81
81
– Now we can add the two half-reactions to obtain the overall cell reaction and cell emf.
V0.40E;e2)aq(Cd)s(Cdo2
====++++→→→→−−−−++++
V0.80E);s(Ag2e2)aq(Ag2o
====→→→→++++−−−−++++
V1.20E);s(Ag2)aq(Cd)aq(Ag2)s(Cdocell
2 ====++++→→→→++++++++++++
– The corresponding cell notation would be
)s(Ag|)M1(Ag||)M1(Cd|)s(Cd2 ++++++++
Do Exercise 19.7 See Problems 19.61-64
82
82
– Note that the emf of the cell equals the standard electrode potential of the cathode minus the standard electrode potential of the anode.
oanode
ocathode
ocell EEE −−−−====
• Calculate the standard emf for the following
voltaic cell at 25°C using standard electrode
potentials. What is the overall reaction?
)s(Fe|)aq(Fe||)aq(Al|)s(Al23 ++++++++
83
83
– The reduction half-reactions and standard potentials are
V1.66E);s(Ale3)aq(Alo3 −−−−====→→→→++++
−−−−++++
V0.41E);s(Fee2)aq(Feo2
−−−−====→→→→++++−−−−++++
V1.66E;e3)aq(Al)s(Al o3 ====++++→→→→−−−−++++
V0.41E);s(Fee2)aq(Feo2
−−−−====→→→→++++−−−−++++
)s(Fe|)aq(Fe||)aq(Al|)s(Al23 ++++++++
– You reverse the first half-reaction and its
half-cell potential to obtain
84
84
– To obtain the overall reaction we must balance the electrons.
V1.66E;e6)aq(Al2)s(Al2o3 ====++++→→→→
−−−−++++
V0.41E);s(Fe3e6)aq(Fe3o2
−−−−====→→→→++++−−−−++++
)s(Fe|)aq(Fe||)aq(Al|)s(Al23 ++++++++
– Now we add the reactions to get the overall cell reaction and cell emf.
V1.66E;e6)aq(Al2)s(Al2o3 ====++++→→→→
−−−−++++
V0.41E);s(Fe3e6)aq(Fe3o2
−−−−====→→→→++++−−−−++++
V25.1E);s(Fe3)aq(Al2)aq(Fe3)s(Al2o32
====++++→→→→++++++++++++
85
85
How do you determine the
direction of spontaneity?
Look at example 19.7 and example 19.8
See problems 19.65 and 19.66
∆G
86
86
Look at Concept check 19.2
Page 790
87
87
Equilibrium Constants from emf’s
• Some of the most important results from
electrochemistry are the relationships among
E°cell, free energy, and equilibrium constant.
– In Chapter 18 we saw that DG equals the maximum
useful work of a reaction.
– For a voltaic cell, work = -nFEo, so when reactants are in their standard states
oonFEG −−−−====∆∆∆∆
– The measurement of cell emf’s gives you yet another way of calculating equilibrium constants.
88
88
– Combining the previous equation, ∆Go = -nFEocell,
with the equation ∆Go = -RTlnK, we get
KlnRTnFEocell
====
KlognF
RT303.2E
ocell
====
– Or, rearranging, we get
– Substituting values for the constants R and F at 25 oC gives the equation
Klogn
0592.0E
ocell
====
(values in volts at 25 oC)
89
89
Look at Examples 19.9 and 19.10
Do Exercises 19.10 and 19.11
Look at problems 19.73 & 74, 81 & 82
Do Exercise 19.9 and see problems 19.69 and 19.70
90
90
– Figure 19.7 summarizes the various relationships among K, ∆Go,
and Eocell.
91
91
A Problem To Consider
• The standard emf for the following cell is 1.10 V.
)s(Cu|)aq(Cu||)aq(Zn|)s(Zn22 ++++++++
Calculate the equilibrium constant Kc for the reaction
)s(Cu)aq(Zn)aq(Cu)s(Zn22
++++++++++++++++
– Note that n=2. Substituting into the equation relating Eo
cell and K gives
Klog2
0592.0V10.1 ====
KlognF
RT303.2E o
cell====
92
92
– Solving for log Kc, you find
37.2Klog ====
– Now take the antilog of both sides:
37c 101.637.2)log(antiK ××××========
– The number of significant figures in the answer equals the number of decimal places in 37.2 (one). Thus
37c 102K ××××====
93
93
Dependence of emf on Concentration
• Recall that the free energy change, ∆G, is related
to the standard free energy change, ∆G°, by the
following equation.
– Here Q is the thermodynamic reaction
quotient.
QlnRTGGo
++++∆∆∆∆====∆∆∆∆
QlnRTnFEnFEocellcell
++++−−−−====−−−−
– Combining the previous equation, ∆Go = -nFEocell,
with the equation ∆Go = -RTlnK, we get
94
94
Dependence of emf on Concentration
• The result rearranges to give the Nernst equation,
an equation relating the cell emf to its standard
emf and the reaction quotient.
QlognF
RT303.2EE
ocellcell
−−−−====
– Substituting values for R and F at 25 oC, we get
Qlogn
0592.0EE
ocellcell
−−−−====
(values in volts at 25 oC)
95
95
Dependence of emf on Concentration
• The result rearranges to give the Nernst equation,
an equation relating the cell emf to its standard
emf and the reaction quotient.
– The Nernst equation illustrates why cell emf
decreases as the cell reaction proceeds.
– As reactant concentrations decrease and
product concentrations increase, Q increases, thus increasing log Q which in turn decreases the cell emf.
96
96
A Problem To Consider
• What is the emf of the following voltaic cell at 25°C?
)s(Cu|)M100.0(Cu||)M101(Zn|)s(Zn252 ++++−−−−++++
××××
The standard emf of the cell is 1.10 V.
– The cell reaction is
)s(Cu)aq(Zn)aq(Cu)s(Zn22
++++++++++++++++
– The number of electrons transferred is 2; hence n = 2. The reaction quotient is
45
2
2
1000.1100.0
1000.1
]Cu[
]Zn[Q
−−−−
−−−−
++++
++++
××××====××××
========
97
97
– The standard emf is 1.10 V, so the Nernst equation becomes
Qlogn
0592.0EE
ocellcell
−−−−====
)1000.1log(2
0592.0V10.1E
4cell
−−−−××××−−−−====
V22.1)12.0(V10.1Ecell====−−−−−−−−====
– The cell emf is 1.22 V.
Do Exercise 19.13 Look at Problems 19.85 &86
98
98
Concept Check 19.3
Page 796
99
99
Some Commercial Voltaic Cells
• The Leclanché dry cell, or zinc-carbon dry cell, is
a voltaic cell with a zinc can as the anode and a
graphite rod in the center surrounded by a paste of
manganese dioxide, ammonium and zinc
chlorides, and carbon black, as the cathode.
– The electrode reactions are
−−−−++++++++→→→→ e2)aq(Zn)s(Zn
2
→→→→++++++++−−−−++++
e2)s(MnO2)aq(NH2 24 )aq(NH2)l(OH)s(OMn 3232++++++++
anode
cathode
(see Figure 19.9)
100
100
Figure 19.9:
Leclanché Dry Cell
101
101
Some Commercial Voltaic Cells
• The alkaline dry cell, is similar to the Leclanché
cell, but it has potassium hydroxide in place of
ammonium chloride.
– The electrode reactions are−−−−++++
++++→→→→ e2)aq(Zn)s(Zn2
→→→→++++++++−−−−
e2)l(OH)s(MnO 22 )aq(OH2)s(OMn 32
−−−−++++
anode
cathode
(see Figure 19.10)
102
102
Figure 19.10: A Small Alkaline Dry Cell
103
103
Some Commercial Voltaic Cells
• The lithium-iodine battery is a solid state battery
in which the anode is lithium metal and the
cathode is an I2 complex.
– The solid state electrodes are separated by a thin crystalline layer of lithium iodide.
(see Figure 19.11)
– Although it produces a low current, it is very reliable and is used to power pacemakers.
104
104
Figure 19.11:
Solid-State
Lithium-Iodine
Battery
105
105
Some Commercial Voltaic Cells
• The lead storage cell (a rechargeable cell)
consists of electrodes of lead alloy grids; one
electrode is packed with a spongy lead to form the
anode, and the other electrode is packed with lead
dioxide to form the cathode.
)l(OH2)s(PbSO 24++++
– The electrode reactions are
−−−−++++−−−−++++++++→→→→++++ e2)aq(H)s(PbSO)aq(HSO)s(Pb 44
→→→→++++++++++++−−−−−−−−++++
e2)aq(HSO)aq(H3)s(PbO 42
anode
cathode
(see Figure 20.12)
106
106
Figure 19.12: Lead Storage Battery
107
107
Some Commercial Voltaic Cells
• The nickel-cadmium cell (nicad cell) consists of
an anode of cadmium and a cathode of hydrated
nickel oxide on nickel; the electrolyte is potassium
hydroxide.
– The electrode reactions are
−−−−−−−−++++→→→→++++ e2)s()OH(Cd)aq(OH2)s(Cd 2
→→→→++++++++−−−−
e)l(OH)s(NiOOH 2)aq(OH)s()OH(Ni 2
−−−−++++
anode
cathode
(see Figure 19.14)
108
108
Figure 19.14: Nicad Storage Batteries
109
109
Some Commercial Voltaic Cells
• A fuel cell is essentially a battery, but differs by
operating with a continuous supply of energetic
reactants, or fuel.
– For a hydrogen-oxygen fuel cell, the electrode reactions are
−−−−−−−−++++→→→→++++ e4)l(OH)aq(OH4)g(H2 22
)aq(OH4e4)l(OH2)g(O 22
−−−−−−−−→→→→++++++++
anode
cathode
(see Figure 19.15)
110
110
Figure 19.15: Hydrogen-Oxygen Fuel Cell
Research
how a fuel
cell works
and what
different
types exist.
To do:
111
111
Figure 19.16: The Electrochemical Process
Involved in the Rusting of Iron
112
112
Figure 19.17:
Cathodic
Protection of
Buried Steel
Pipe
113
113
Figure 19.18: A Demonstration of
Cathodic Protection
Unprotected nail
Fe2+ + Ferricyanide →
114
114
115
115
Electrolytic Cells
• An electrolytic cell is an electrochemical cell in
which an electric current drives an otherwise
nonspontaneous reaction. (See Video: Electrolysis
of Water)
– The process of producing a chemical change in an electrolytic cell is called electrolysis.
– Many important substances, such as aluminum metal and chlorine gas are produced commercially by electrolysis.
116
116
Figure 19.19: Electrolysis of Molten
Sodium Chloride
117
117
Electrolysis of Molten Salts
• A Downs cell is a commercial electrochemical cell
used to obtain sodium metal by electrolysis of
molten NaCl. (see Figure 19.20)
– A number of other reactive metals are obtained by the electrolysis of a molten salt.
– Lithium, magnesium, and calcium metals are all obtained by the electrolysis of the chlorides.
Do Exercise 19.15 See Problems 19.91-92
118
118
Figure 19.20: Downs Cell for Preparation of
Sodium Metal
119
119
Figure 19.22: Chlor-Alkali Membrane Cell
120
120
Figure 19.23: Chlor-Alkali Mercury Cell
121
121
Figure 19.24: Purification of Copper by
Electrolysis
122
122
Figure 19.24: Purification of Copper by
Electrolysis (cont’d)
123
123
124
124
Stoichiometry of Electrolysis
• What is new in this type of stoichiometric
problem is the measurement of numbers of
electrons.
– You do not weigh them as you do substances.
– Rather, you measure the quantity of electric chargethat has passed through a circuit.
– To determine this we must know the current and the length of time it has been flowing.
125
125
– Electric current is measured in amperes.
– An ampere (A) is the base SI unit of current equivalent to 1 coulomb/second.
– The quantity of electric charge passing through a circuit in a given amount of time is given by
Electric charge(coul) = electric current (coul/sec) ×××× time lapse(sec)
126
126
A Problem To Consider
• When an aqueous solution of potassium iodide is
electrolyzed using platinum electrodes, the half-
reactions are
How many grams of iodine are produced when a current of
8.52 mA flows through the cell for 10.0 min?
−−−−−−−−++++→→→→ e2)aq(I)aq(I2 2
)aq(OH2)g(He2)l(OH2 22
−−−−−−−−++++→→→→++++
– When the current flows for 6.00 x 102 s (10.0 min), the amount of charge is
C11.5)s1000.6()A1052.8(23
====××××××××××××−−−−
127
127
Figure 19.21: Electrolysis of Aqueous
Potassium Iodide
128
128
A Problem To Consider
How many grams of iodine are produced when a current of
8.52 mA flows through the cell for 10.0 min?
−−−−−−−−++++→→→→ e2)aq(I)aq(I2 2
)aq(OH2)g(He2)l(OH2 22
−−−−−−−−++++→→→→++++
– Note that two moles of electrons are equivalent to one mole of I2. Hence,
−−−−
−−−−
××××××××
××××
emol2
Imol1
C1065.9
emol1C11.5 2
4 23
2
2 Ig1073.6Imol1
Ig254 −−−−××××====××××
• When an aqueous solution of potassium iodide is
electrolyzed using platinum electrodes, the half-
reactions are
129
129
Do Exercise 19.16
See Problems 19.93-94
Do Exercise 19.17
See Problems 19.95-96
Do Exercise 19.18
See problems 19.97-98
130
130
Operational Skills
• Balancing oxidation-reduction reactions
• Sketching and labeling a voltaic cell
• Writing the cell reaction from the cell notation
• Calculating the quantity of work from a given amount of cell reactant
• Determining the relative strengths of oxidizing and reducing agents
• Determining the direction of spontaneity from electrode potentials
• Calculating the emf from standard potentials
131
131
Operational Skills
• Calculating the free-energy change from electrode
potentials
• Calculating the cell emf from free-energy change
• Calculating the equilibrium constant from cell emf
• Calculating the cell emf for nonstandard conditions
• Predicting the half-reactions in an aqueous
electrolysis
• Relating the amounts of product and charge in an
electrolysis
132
132
End of Chapter 19
Related Documents
