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Chemistry The Final Conflict Chapter 18: Electrochemistry and Electric Vehicles

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Title: Chemistry The Final Conflict Chapter 18: Electrochemistry and Electric Vehicles


1
ChemistryThe Final ConflictChapter 18
Electrochemistryand Electric Vehicles
2
Electrochemistry The area of chemistry that
examines the transformations between chemical and
electrical energy. Oxidation-reduction (redox)
reactions involve an exchange of electrons
between reacting species (see chapter 4,
section 4.8). You may need to review the
following terminology oxidation, reduction,
oxidizing agent, reducing agent and half-reaction
. In the following reaction, what is being
oxidized, reduced? What is the oxidizing agent,
the reducing agent?
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Fig. 18.1 A piece of zinc is immersed in a
copper(II) sulfate solution. The Cu(II) is
spontaneously converted to elemental Cu and the
solid Zn dissolves as Zn2 ions.
Cu2(aq) Zn(s) ? Cu(s) Zn2(aq)
What are the half-reactions? What is oxidized?
reduced? In this example, there is an exchange of
electrons between the oxidized and reduced
species that is thermodynamically favored
(exergonic). The goal of using an
electrochemical cell is to extract usable work
from this electron transfer.
5
Problem 12. Identify which elements (if any)
undergo oxidation or reduction. 4 ClO3-(aq) ?
Cl- 3 ClO4-
6
Problem. Sometimes the cell reaction of NiCd
batteries is written with Cd metal as the anode
and solid NiO2 as the cathode. Assuming the
products of the electrode reactions are solid
hydroxides of Cd(II) and Ni(II), respectively,
write a balanced chemical equation for the cell
reaction.
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Fig. 18.2 An electrochemical cell is a reaction
system in which oxidation and reduction reactions
occur in separate compartments (or cells) either
consume or produce electrical energy. The cells
are separated by a salt bridge or semi-permeable
membrane that allows ions to migrate from one
cell to the other. Electrons move from anode
(oxd) to cathode (red).
Voltaic or Galvanic cell chemical energy is
used to produce electrical energy (DGlt0) (i.e. a
battery). Electrolytic cell an external source
of electrical energy is used to do work on a
chemical system (i.e. charging a car battery with
the alternator after the car has started).
9
Alessandro Volta is credited with building the
first batterywhich was built by alternating
layer of zinc and silver with paper-soaked brine
between the metals. Volta coined the term
galvanism to distinguish the animal
electricity
observed by his adversary, Luigi Galvani
10
Anode oxidation
Cathode reduction
AgNO3(aq)
Ni(NO3)2(aq)
Write the balanced redox reaction for this
electrochemical cell. Which direction will the
nitrate ions flow in the salt bridge? (see sample
exercise 18.2)
11
Anode oxidation
Cathode reduction
AgNO3(aq)
Ni(NO3)2(aq)
12
Cell potential (Ecell) or Electromotive Force
(emf) is the voltage between the electrodes of a
voltaic cell. A Faraday (F) is the electrical
charge on one mole of electrons or 9.65E4
Coulombs (C)/ mol e?. (The charge on one
electron is -1.602E-19 C). The quantity of charge
flowing through an electrical circuit is C
n?F Electrical work is done when a charge moves
through an electrical potential, welec
?C?Ecell n?F?Ecell Since free energy is that
available to do work. this equation
becomes DG n?F?Ecell 1 C?V 1 Joule
(J)
13
Fuel cells use a controlled electron transfer
between hydrogen and oxygen to produce electrical
energy. H2(g) ½ O2(g) ? H2O(l) DG
?237 kJ What is the electromotive potential that
can be produced by this cell under standard
conditions (Eo)?
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A standard potential (Eo) is the electromotive
force of a half-reaction written as a reduction
reaction in which all reactants and products are
in their standard states (see Table A5.4 or
en.wikipedia.org/wiki/Table_of_standard_electrode_
potentials). The standard cell potential
(Eºcell) is the potential of a cell when all
reactants and products are in their standard
states, i.e. the pressure of all gases are 1 bar,
and the concentration of dissolved species are 1
molar. Eºcell Eºcathode Eºanode or
Eºcell Eºreduction Eºoxidation
17
Fig. 17.7 The Standard Hydrogen Electrode (SHE)
has a solution of 1 M HCl and is bathed in a
stream of H2 gas at 1 bar (1 atm) pressure. This
half-reaction has a defined potential of 0.00 V
as either a reduction or oxidation reaction and
is used to reference the potentials of other
half-reactions.
18
Also see Table A6.1 in your text. This source
http//www.jesuitnola.org/upload/clark/refs/red_po
t.htm
19
Fig. 17.8 The SHE can be used in a cell with
either oxidation or reduction reactions to
determine the standard potential of that
half-reaction. Use the above figures and the
following cell potential equation to find the
standard potentials for the Cu and Zn
reactions. Eºcell Eºcathode Eºanode
20
SHE 2H(aq) 2e- ? H2(g) E0 0.0000 Zn2(aq)
2e- ? Zn(s) E0 -0.7618 Cu2(aq) 2e- ?
Cu(s) E0 0.52
21
One of the first batteries built was that of
Allesandro Volta (1745-1827). Calculate the
standard cell potential (Eºcell ) for this
battery that had a Ag/Ag cell connected to a
Zn/Zn2 cell by a salt bridge (a blotter soaked
with a NaNO3 solution). Eºcell Eºcathode
Eºanode From Table A6.1 Ag 1 e? ? Ag(s)
E? 0.7996 V Zn2 2 e? ? Zn(s) E?
-0.7618 V Compare two reduction half-reactions
the more positive of the two will occur as the
cathode reaction (for a voltaic cell). The other
reaction will be the anode.
See sample and practice exercise 18.4.
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Problem. A Voltaic cell based on the following
pair of half-reactions is constructed. Write a
balanced equation for the overall cell reaction,
and identify which half-reaction takes place at
the anode and cathode. AgBr(s) e- ? Ag(s)
Br-(aq) E 0.095V MnO2(s) 4 H 2e- ?
Mn2(aq) 2 H2O(l) E 1.23V
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The Nernst Equation (Walther Nernst, 1864-1941)
can be used to calculate the cell potential for
non-standard conditions usually when the conc. of
dissolved species ? 1 . Where R is the gas
constant, T is the temperature in Kelvin, n is
the number of moles of electrons transferred
between the oxidation and reduction reactions, F
is the Faraday constant and Q is the reaction
quotient for the system. What is the value of Q
when all the reaction species are at standard
conditions?
Nernst Eq. at 25ºC.
26
Fig. 18.5 The standard lead-acid storage battery
found in your car uses the following
reaction. Pb(s) PbO2(s) 2 H2SO4(aq) ? 2
PbSO4(s) 2 H2O(l) With an electrolyte of 4.5 M
H2SO4 each cell produces 2.0 V.
27
Fig. 17.10 The cell potential will decrease as
reactants are converted to products. This is an
example of the cell potential in a standard
lead-acid car battery as the sulfuric acid is
used up while the battery discharges.
Pb(s) PbO2(s) 2 H2SO4(aq) ? 2 PbSO4(s) 2
H2O(l)
28
Problem 18.53. Permanganate ion can oxidize
sulfite to sulfate in basic solution as
follows 2MnO4(aq) 3SO32(aq) H2O( ) ?
2MnO2(s) 3SO42(aq) 2OH(aq) Determin
e the Standard Potential for the reaction at 298
K and when the concentrations of the reactants
and products are as follows MnO4 0.150 M,
SO32 0.256 M, SO42 0.178 M, and OH
0.0100 M. Will the value of Er?n increase or
decrease as the reaction proceeds?
29
Also see Table A6.1 in your text. This source
http//www.jesuitnola.org/upload/clark/refs/red_po
t.htm
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The Nernst Equation can be used to predict the
cell potentials under non-standard conditions.
In this example (A) the two identical cells with
differing concentrations of dissolved silver will
produce a cell potential (see p. 867).
33
At equilibrium the Nernst Equation changes
since at equilibrium Ecell 0 and Q K, so the
equation can be rearranged as follows
34
Using standard reduction potentials from the
table in Appendix , calculate the value of the
equilibrium constant for the following reaction
at 298 K. 5 Fe2(aq) MnO4(aq) 8 H(aq)
? 5 Fe3(aq) Mn2(aq) 4 H2O(l)
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Fig. 17.1 the predicted effect of temperature on
the cell potential of a lead-acid battery using
the Nernst Eqn.
Other T-dependent factors have a greater
influence on battery performance.
38
Differences between an electrolytic cell and a
voltaic cell.
39
From a practical stand point one of the important
characteristics of a battery is its ability to do
work. wcell CEcell Battery capacity can be
expressed in coulombs x volts ( joules). Other
common definitions of battery capacity are
useful, for example 1 ampere (amp) 1
coulomb(C)/sec or 1 C 1 ampsec The Faraday
constant can be written F 9.65E4 A?s/mole 1
watt 1 voltamp 1 J/s The watt (James Watt,
1736-1819) is a widely used unit electrical
power. Consequently, a cell producing 1 volt of
potential and 1 amp of current will produce 1
watt of power.
40
Problem 78. In the electrolysis of water, how
long will it take to produce 100.L of H2(g) at
STP using an electrolysis cell through which
flows a current of 50.mA?
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42
Fig. 17.13 The discharge-charging cycle of a
lead-acid battery.
43
Problems. If it takes 6.0 seconds of discharge
for a car battery to start the engine and the
starter drew a current of 230 A, what mass of Pb
will be oxidized to PbSO4 in this time? How
long will it take to re-charge the battery with
an alternator current of 30.0 A?
Pb(s) PbO2(s) 2 H2SO4(aq) ? 2 PbSO4(s) 2
H2O(l)
See sample practice exercise 18.7 and assume
100 efficiency in these reactions.
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a. What will be the oxidation reaction that
occurs in the above voltaic cell? b. What will
be the reduction reaction that occurs in the
above cell? c. What will be the standard cell
potential for the above cell at 25 ?C? d. On the
diagram, indicate the direction of electron flow
through the external wire. e. On the
diagram, indicate the anode and cathode
compartments. f. On the diagram, indicate the
direction of migration of sulfate through the
salt bridge? g. Calculate the cell potential
when the concentration of Cd2 is 0.050 M and the
concentration of Mg2 is 0.0025 M (at 298 K).
46
Fig. 17.15 Thin coatings of metals can be
applied using electrolytic reactions. How many
grams of silver can be plated from a silver
nitrate solution using a 20. mA current for 15
minutes?
47
Fuel cells are a voltaic device in which there is
a flow of reactants to the anode and
cathode. The hydrogen fuel cell uses streams of
H2 and O2 gases that diffuse to the following
anode and cathode reactions, respectively. H2(g)
? 2 H(aq) 2 e Eº 0.000 V O2(g) 4
H(aq) 4 e ? 2 H2O(l) Eº 1.229 V The
overall reaction is 2 H2(g) O2(g) ? 2
H2O(l) Eºcell 1.229 V
48
17_04.jpg
Fig. 17.4 The disposable zinc-acid (dry cell)
battery.
Zn 2 NH4Cl 2 MnO2 ? 2 Zn(NH3)2Cl2 Mn2O3
H2O, Ecell1.5 V
49
17_05.jpg
Fig. 17.5 The alkaline cell has the same
potential as the classic dry cell but the zinc
anode is oxidized to Zn(OH)2.
50
17_06.jpg
Fig. 17.6 The Ni-Cd or nickel-cadmium battery
can be recharged because the products adhere to
the respective electrodes and the reaction can be
readily reversed.
Cd(s) NiO(OH)(s) 2 H2O(l) ? Cd(OH)2(s) 2
Ni(OH)2(s)
51
17_09.jpg
Fig. 17.9 In the nickel-metal hydride battery
(NiMH) , hydrogen atoms are stored in the
interstitial spaces of the metal matrix. They
can migrate from these spaces to participate in
the anode half-reaction. NiO(OH) is
simultaneously reduced at the cathode.
52
In the lithium-ion battery, Li is stored in pure
graphite anodes (see section 10.8 each six
carbon ring of graphite stores 1 lithium ion).
The cathodes are made of porous transition metal
oxides ( i.e. MnO2) which can form highly stable
complexes with Li ions. In a fully charge
battery there is a concentration gradient between
the anode and the cathode. The lithium ions
migrate down the gradient and at the same time
electrons flow in the external circuit to balance
the charge. The electrodes in the Li-ion cell
react with oxygen and water so they must be
either entirely solid-state or use non-aqueous
solvents. See pages 862-863 for description of
the Li-ion cell.
53
Some common types of batteries and uses.
54
Cell Potential Tutorial
You should view these tutorials on your own.
They are found on your CD or the publishers
website.
PC version Mac version
This tutorial explores the concept of cell
potential (Ecell) as a measure of how much
electrical energy is stored in an electrochemical
cell. Includes practice exercises.
55
17_16.jpg
Fig. 17.16
56
17_18.jpg
Fig. 17.18 Photochemical cells can be used to
directly convert sunlight energy to electricity.
This cells uses sunlight to catalyze the
reduction of water to H2 at the cathode and
oxidation to form O2 at the anode.
57
17_19.jpg
Fig. 17.19 Biological batteries can use
bacteria to catalyze redox reactions and harness
the electron flow to do useful work.
58
Fuel Cell Tutorial
You should view these tutorials on your own.
They are found on your CD or the publishers
website.
PC version Mac version
Learn how fuel cells use a redox reaction between
hydrogen and oxygen to produce electrical energy.
Includes practice exercises.
59
W. W. Norton CompanyIndependent and
Employee-Owned
  • This concludes the Norton Media Libraryslide
    set for chapter 17ChemistryThe Science in
    Context byThomas Gilbert,Rein V. Kirss,
    Geoffrey Davies

60
Zinc-Copper Cell Tutorial
You should view these tutorials on your own.
They are found on your CD or the publishers
website.
PC version Mac version
This tutorial illustrates the reactions that
occur at the electrodes of a typical zinc copper
battery, and explores how the energy released by
a voltatic cell is used to do work on the
surroundings. Includes practice exercises.
61
Free Energy Tutorial
You should view these tutorials on your own.
They are found on your CD or the publishers
website.
PC version Mac version
Learn how the potential of an electrochemical
cell can be used to determine the free energy
available to do work, and explore the
relationships between free energy, cell potential
and equilibrium constant. Includes practice
exercises.
62
Ni-Cd Battery Tutorial
You should view these tutorials on your own.
They are found on your CD or the publishers
website.
PC version Mac version
This tutorial explores the oxidation-reduction
reactions that power rechargeable Ni-Cd batteries
and describes the changes in reaction quotient as
a battery loses its charge or is recharged.
Includes practice exercises.
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