Transcript Document

Chemistry
The Final Conflict…
Chapter 18: Electrochemistry
and Electric Vehicles
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?
2
2 FeS2 (s, pyrite) 7 O2 (g)  2H2O(l) 2 Fe2 (aq)  4 SO 4 (aq)  4 H (aq)
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.
Problem 12. Identify which elements (if any)
undergo oxidation or reduction.
4 ClO3-(aq) →
Cl- + 3 ClO4-
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.
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 semipermeable 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 (DG<0) (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).
Alessandro Volta is credited
with building the first
battery…which 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
Anode =
oxidation
Ni(NO3)2(aq)
Cathode = reduction
AgNO3(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)
Anode =
oxidation
Ni(NO3)2(aq)
Cathode = reduction
AgNO3(aq)
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 C۰V = 1 Joule (J)]
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)?
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_poten
tials).
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
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.
Also see Table A6.1 in your text.
This source:
http://www.jesuitnola.org/upload/clark/refs/red_pot.htm
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
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
One of the first batteries built was that of Allesandro Volta (17451827). 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.
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 halfreaction 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
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 .
RT
Ecell  E cell 
 lnQ
nF
o
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?
0.0592 V
Ecell  E cell 
 log Q
n
o
Nernst Eq. at 25ºC.
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.
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)
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)
Determine 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?
Also see Table A6.1 in your text.
This source:
http://www.jesuitnola.org/upload/clark/refs/red_pot.htm
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).
At equilibrium the Nernst Equation:
0.0592
Ecell  E cell 
 log Q
n
o
changes since at equilibrium Ecell = 0 and Q = K, so
the equation can be rearranged as follows:
0.0592
0  E cell 
 log K
n
o
nE cell
log K 
0.0592
o
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)
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.
Differences between an electrolytic cell and a voltaic cell.
From a practical stand point one of the important
characteristics of a battery is its ability to do work.
wcell = C·Ecell
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 amp·sec
The Faraday constant can be written: F = 9.65E4 A۰s/mole
1 watt = 1 volt·amp = 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.
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?
Fig. 17.13: The discharge-charging cycle of a lead-acid battery.
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?
Pb(s) + PbO2(s) + 2 H2SO4(aq)  2 PbSO4(s) + 2 H2O(l)
How long will it take to re-charge the battery with an
alternator current of 30.0 A?
See sample & practice exercise 18.7 and assume 100% efficiency in
these reactions.
Mg
MgSO4(1 M, aq)
a.
b.
c.
d.
Cd
Salt
bridge
CdSO4(1 M, aq)
What will be the oxidation reaction that occurs in the above voltaic cell?
What will be the reduction reaction that occurs in the above cell?
What will be the standard cell potential for the above cell at 25 C?
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).
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?
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
Fig. 17.4: The
“disposable” zincacid (dry cell)
battery.
Zn + 2 NH4Cl + 2 MnO2  2 Zn(NH3)2Cl2 + Mn2O3 + H2O, Ecell=1.5 V
17_04.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.
17_05.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)
17_06.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.
17_09.jpg
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.
Some common types of batteries and uses.
You should view these tutorials on your own. They are
found on your CD or the publishers website.
Cell Potential Tutorial
»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.
Fig. 17.16:
17_16.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.
17_18.jpg
Fig. 17.19: “Biological” batteries can use bacteria to catalyze redox
reactions and harness the electron flow to do useful work.
17_19.jpg
You should view these tutorials on your own. They are
found on your CD or the publishers website.
Fuel Cell Tutorial
»PC version
»Mac version
Learn how fuel cells use a redox reaction between
hydrogen and oxygen to produce electrical energy.
Includes practice exercises.
W. W. Norton & Company
Independent and Employee-Owned
This concludes the Norton Media Library
slide set for chapter 17
Chemistry
The Science in Context
by
Thomas Gilbert,
Rein V. Kirss, &
Geoffrey Davies
You should view these tutorials on your own. They are
found on your CD or the publishers website.
Zinc-Copper Cell Tutorial
»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.
You should view these tutorials on your own. They are
found on your CD or the publishers website.
Free Energy Tutorial
»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.
You should view these tutorials on your own. They are
found on your CD or the publishers website.
Ni-Cd Battery Tutorial
»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.