Chapter 21 Electrochemistry

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Chapter 21Electrochemistry
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Electrochemical Processes
Chemical processes can either release energy or
absorb energy. The energy can sometimes be in
the form of electricity. Electrochemistry has
many applications in the home as well as in
industry. Flashlight and automobile batteries
are examples of devices used to generate
electricity. Biological systems also use
electrochemistry to carry nerve impulses.
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Spontaneous Redox Reaction
When a strip of zinc metal is dipped into an
aqueous solution of blue copper sulfate, the zinc
becomes copper-plated. Zn(s) Cu2(aq)
Zn2 (aq) Cu(s) The net ionic equation
involves only zinc and copper. Electrons are
transferred from zinc atoms to copper ions. This
is a redox reaction that occurs spontaneously
4
Spontaneous Redox Reaction
As the reactions proceeds, zinc atoms lose
electrons as they are oxidized to zinc ions. The
zinc metal slowly dissolves. 0
2 2
0 Zn(s) Cu2(aq)
Zn2 (aq) Cu(s) At the same time, copper ions
in solution gain the electrons lost by the zinc.
They are reduced to copper atoms are deposited
as metallic copper.
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Spontaneous Redox Reaction
As the copper ions in solution are gradually
replace by zinc ions, the blue color of the
solution fades. Oxidation Zn(s)
Zn2 (aq) 2e- Reduction Cu2(aq) 2e-
Cu(s)
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Activity Series of Metals
Zinc is higher on the list than copper. For any
two metals in an activity series, the more active
metal is the more readily oxidized.
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Electric Current
When zinc is dipped into a copper sulfate
solution, zinc becomes plated with copper. In
contrast, when a copper strip is dipped into a
solution of zinc sulfate, the copper does not
spontaneously become zinc-plated. This is
because copper metal is not oxidized by zinc
ions. When a zinc strip is dipped into a copper
sulfate solution, electrons are transferred from
zinc metal to copper ions. This flow of
electrons is an electric current.
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Electrochemical Process
The zinc-metalcopper-ion system is an example of
the conversion of chemical energy into electrical
energy. Electrochemical process any conversion
between chemical energy and electrical energy All
electrochemical processes involve redox
reactions. If a redox reaction is to be used a a
source of electrical energy, the two
half-reactions must be physically separated.
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Electrochemical Cell
The electrons released by zinc must pass through
an external circuit to reach the copper ions if
useful electrical energy is to be produced. In
this case the system serves as an electrochemical
cell. Also, an electric current can be used to
produce a chemical change. That system, too,
serves as an electrochemical cell.
Electrochemical cell any device that converts
chemical energy into electrical energy or vice
versa.
10
Voltaic Cells
In 1800, Italian physicist Alessandro Volta build
the first electrochemical cell that could be used
to generate a dire electric current. Voltaic
cells are electrochemical cells used to convert
chemical energy into electrical energy.
Electrical energy is produced in a voltaic cell
by spontaneous redox reactions within the cell.

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Constructing a Voltaic Cell
Half cell one part of a voltaic cell in which
either oxidation or reduction occurs. Typical
half-cell consists of a piece of metal immersed
in a solution of its ions. Example one
half-cell is a zinc rod immersed in a solution of
zinc sulfate. Other half-cell is a copper rod
immersed in a solution of copper sulfate. Half
cells are connected by a salt bridge which is a
tube containing a strong electrolyte, often
potassium sulfate.
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Constructing a Voltaic Cell
A porous plate my be used instead of a salt
bridge The porous plate allows ions to pass
from on half-cell to the other but prevents the
solutions from mixing completely. A wire
carries the electrons in the external circuit
from the zinc rod to the copper rod. A voltmeter
or light bulb can be connected in the circuit.
The driving force of such a voltaic cell is the
spontaneous redox reaction between zinc metal and
copper(II) ions in solution.
13
Constructing a Voltaic Cell
The zinc and copper rods in this voltaic cell are
the electrodes. Electrode a conductor in a
circuit that carries electrons to or from a
substance other than a metal. The reaction
at the electrode determines whether the electrode
is labeled as an anode or a cathode Anode the
electrode at which oxidation occurs Cathode the
electrode at which reduction occurs.
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Constructing a Voltaic Cell
Electrons are consumed at the cathode and its
labeled the positive electrode. (reduction
occurs) Electrons are produced at the anode and
its labeled the negative electrode. (oxidation
occurs) The reaction at the electrode
determines whether the electrode is labeled as an
anode or a cathode All parts of the voltaic cell
remain balanced in terms of charge at all times.
The moving electrons balance any charge that
might build up as oxidation and reduction occur.

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Questions
At which electrode does oxidation take place? At
the anode (negative electrode) Where does
reduction take place? At the cathode (positive
electrode) What path do the electrons given up by
zinc follow? They go through the wire and the
electric light to the copper electrode. What
happen to the electrons at the copper
electrode? They reduce copper ions to copper.
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How a Voltaic Cell Works

• These steps actually occur at the same time.
• Zn(s) Zn 2(aq) 2e-
• Electrons are produced at the zinc rod according
  to the oxidation half reaction
• The electrons leave the zinc anode and pass
  through the external circuit to the copper rod.
• Electrons enter the copper rod and interact with
  copper ions in solution.
• Cu 2 (aq) 2e- Cu(s)

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How a Voltaic Cell Works

• To complete the circuit, both positive and
  negative ions move through the aqueous solutions
  via the salt bridge.
• The overall cell reaction (note the electrons in
  the overall reaction must cancel.
• Zn(s) Zn2(aq) 2e-
• Cu2 (aq) 2e- Cu(s)
• Zn(s) Cu2 (aq) Zn2(aq)
  Cu(s)

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The Need for a Salt Bridge
As zinc is oxidized at the anode, Zn2 ions enter
the solution and they have no negative ions to
balance their charges. So a positive charge tends
to build up around the anode. At the cathode,
Cu2 ions are reduced to Cu and taken out of the
solution leaving behind unbalanced negative ions.
Thus, a negative charge tends to develop around
the cathode.
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The Need for a Salt Bridge
The salt bridge allows negative ions, such as
SO42-, to be drawn to the anode to balance the
growing positive charge. Positive ions, such as
K, are drawn from the salt bridge to balance the
growing negative charge at the cathode.
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Representing Electrochemical Cells
You can represent the zinc-copper voltaic cell by
suing the following shorthand form. Zn(s)
ZnSO4(aq) CuSO4(aq) Cu(s) The single vertical
lines indicate boundaries of phases that are in
contact. The double vertical lines represent the
salt bridge or porous partition that separates
the anode compartment from the cathode
compartment. The half-cell that undergoes
oxidation (the anode) is written first.
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Using Voltaic Cells as Energy Sources
The zinc-copper voltaic cell is no longer used
commercially. Current technologies that use
electrochemical processes to produce electrical
energy include dry cells, lead storage batteries
and fuel cells. Dry Cells a voltaic cell in
which the electrolyte is a paste. Dry cells used
when a compact, portable electrical energy source
is required.
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Dry Cells
A common type of dry cell is a flashlight
battery. A zinc container is filled with a thick,
moist electrolyte paste of manganese (IV) oxide,
zinc chloride, ammonium chloride, and water. A
graphite rode is embedded in the paste. The
zinc container is the anode and the graphite rod
is the cathode. The thick paste and its
surrounding paper liner prevent the contents of
the cell from freely mixing, so a salt bridge is
not needed.
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Dry Cells
Oxidation Zn (s) Zn2 (aq)
2e- Reduction 2MnO2(s) 2NH4(aq) 2e-
Mn2O3(s) 2NH3(aq)
H2O(l)
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Dry Cells
In an ordinary dry cell, the graphite rod serves
only as a conductor and does not undergo
reduction. MnO2 is the species that is actually
reduced. The electrical potential of this cell
starts out at 1.5V but decreases steadily during
use to about 0.8V. Dry cells of this type are
not rechargeable because the cathode reaction is
not reversible.
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Alkaline Battery
The alkaline battery is an improved dry cell used
for the same purposes. In the alkaline battery,
the reactions are similar to those in the common
dry cell, but the electrolyte s a basic KOH past.
This change in design eliminates the buildup of
ammonia gas and maintains the Zn electrode, which
corrodes more slowly under alkaline conditions.
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Alkaline Battery
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Lead Storage Batteries
People depend on lead storage batteries to start
their cars. Battery is a group of cells
connected together. A 12-V car battery consists
of six voltaic cells connected together. Each
cell produces about 2 V and consists of lead
grids.
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Lead Storage Batteries
One set of grids, the anode, is packed with
spongy lead. The other set, the cathode, is
packed with lead(IV) oxide. The electrolyte for
both half-cells in a lead storage batter is
concentrated sulfuric acid. Using the same
electrolyte for both half-cells allows the cell
to operate without a salt bridge or porous
separator.
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Lead Storage Batteries
Pb(s) SO42-(aq) PbSO4(s)
2e- PbO2(s) 4H(aq) SO42-(aq) 2e-
PbSO4(s) 2H2O(l) When a lead storage battery
discharges, it produces the electrical energy
needed to start a car. The overall spontaneous
redox reaction that occurs is the sum of the
oxidation and reduction half-reactions. Pb(s)
PbO2(s) 2H2SO4 (aq) 2PbSO4(s)
2H2O(l) The equation shows that lead sulfate
forms during discharge.
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Lead Storage Batteries
Pb(s) PbO2(s) 2H2SO4 (aq)
2PbSO4(s) 2H2O(l) The sulfate slowly builds up
on the plates, and the concentration of sulfuric
acid decreases. 2PbSO4(s) 2H2O(l)
Pb(s) PbO2(s) 2H2SO4 (aq) The reverse
reaction occurs when a lead storage battery is
recharged. This occurs when the cars generator
is working properly. The reverse reaction is
not a spontaneous reaction. A direct current must
pass through the cell in a direction opposite
that of the current flow during discharge.
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Lead Storage Batteries
In theory, a lead storage battery can be
discharged and recharged indefinitely, but in
practice its lifespan is limited. Small amounts
of lead sulfate fall from the electrodes and
collect on the bottom of the cell. Eventually,
the electrodes lose so much lead sulfate that the
recharging process is ineffective or the cell is
shorted out.
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Lead Storage Battery
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Fuel Cells
To overcome the disadvantages associated with
lead storage batteries, cells with renewable
electrodes have been developed. Fuel cells are
voltaic cells in which a fuel substance undergoes
oxidation and from which electrical energy is
continuously obtained. Fuel cells do not have to
be recharged. They can be designed to emit no
air pollutants and to operate more quietly and
more const-effectively than a conventional
electrical generator
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Fuel Cells
Simplest fuel cell is the hydrogen-0xygen fuel
cell. Fuel cells are voltaic cells in which a
fuel substance undergoes oxidation and from which
electrical energy is continuously obtained. Fuel
cells do not have to be recharged. They can be
designed to emit no air pollutants and to operate
more quietly and more const-effectively than a
conventional electrical generator There are
three compartment separated from one another by
two electrodes made of porous carbon.
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Fuel Cells
Oxygen (the oxidizer) is fed into the cathode
compartment Hydrogen (the fuel) is fed into the
anode compartment. The gases diffuse slowly
through the electrodes. The electrolyte in the
central compartment is a hot, concentrated
solution of potassium hydroxide. Electrons from
the oxidation half-reaction at the anode pass
through an external circuit to enter the
reduction half reaction at the cathode.
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Fuel Cells
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Fuel Cells
Oxidation 2H2(g) 4OH-(aq) 4H2O(l)
4e- Reduction O2(g) 2H2O(l) 4e- 4
OH-(aq) The overall reaction in the
hydrogen-oxygen fuel cell is the oxidation of
hydrogen to form water. 2H2 (g) O2 (g)
2H2O (l) Fuel cells were developed for space
travel where lightweight, reliable power systems
are needed. Fuel cells different from lead
storage batteries in that they are not
self-contained.
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Fuel Cells
Operation depends on a steady flow of fuel and
oxygen into the cell (where combustion takes
place) and the flow of the combustion product out
of the cell. In the case of the hydrogen fuel
cell, the product is pure water. Both the
electricity generated and the water produced are
consumed in space flights. Fuel cells convert
75 of the available energy into electricity.
Conventional electric power plant converts from
35 to 40 of the energy of coal to electricity.
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Fuel Cells
Other fuels, such as methane (CH4) and ammonia
(NH3), can be used in place of hydrogen.
4NH3(g) 3O2 (g) 2N2 (g) 6H2O (g)
CH4(g) 2O2(g) CO2(g) 2H2O(g)
Other oxidizers, such as chlorine (Cl2) and
ozone (O3)can be used in place of oxygen.
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Question
Make a sketch of a tin/lead voltaic cell Sn
SnSO4 PbSO4 Pb Label the cathode and anode,
and indicate the direction of electron flow.
Write the equations for the half-reactions. Sn(s)
Sn2(aq) 2e- Pb2(aq) 2e-
Pb(s) Tin is the anode, lead is the cathode.
The electrons flow from tin to lead.
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Questions
What type of reaction occurs during an
electrochemical process? Redox What is the source
of electrical energy produced in a voltaic
cell? Spontaneous redox reactions within the
cell If the relative activities of two metals are
known, which metal is more easily oxidized? The
metal with the higher activity.
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Homework

• Draw detailed pictures of the following
• A Voltaic Cell
• A Lead Storage Battery
• Fuel Cell
• Make sure to include all substances involved as
  well as the half reactions and overall reaction.
• Make sure you label the anode and cathode and
  indicate where oxidation and reduction take
  place.
• List common uses of each
• List advantages of disadvantages of each.

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End of Section 20.1
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Electrical Potential
Electrical potential of a voltaic cell is a
measure of the cells ability to produce an
electric current. (in Volts) You cannot measure
the electrical potential of an isolated
half-cell. When two half-cells are connected to
form a voltaic cell, the difference in potential
can be measured. The electrical potential of a
cell results from a competition for electrons
between two half-cells. The half-cell that has a
greater tendency to acquire electrons will be the
one in which reduction occurs.
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Electrical Potential
Oxidation occurs in the other half-cell as there
is a loss of electrons. Reduction potential
the tendency of a given half-reaction to occur as
a reduction. The half-cell in which reduction
occurs has a greater reduction potential than the
half-cell in which oxidation occurs. Cell
potential the difference between the reduction
potentials of the two half-cells.
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Cell Potential
reduction potential
reduction potential Cell Potential of
half-cell in which - of half-cell in which
reduction occurs
oxidation occurs E0cell E0red -
E0oxid Standard cell potential (E0cell) is the
measured cell potential when the ion
concentration in the half-cells are 1M, any gases
are at a pressures of 101 kPa, ant the
temperature is 25ºC. The symbols E0red and
E0oxid represent the standard reduction
potential for the reduction and oxidation
half-cells, respectively.
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The Lemon Battery
A working voltaic cell made using a lemon and
strips of copper and zinc. Which is the anode
and which is the cathode? Zn is anode Cu is
cathode What process goes on at the anode and
cathode? Oxidation at anode, reduction at
cathode. Which process are electrons
lost? Oxidation What role does the lemon play in
the battery? Salt bridge
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Standard Hydrogen Electrode
Because half-cell potentials cannot be measured
directly, scientists have chosen an arbitrary
electrode to serve as a reference. The standard
hydrogen electrode is used with other electrodes
so the reduction potentials of the other cells
can be measured. The standard reduction potential
of the hydrogen electrode has been assigned a
value of 0.00 V.
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Standard Hydrogen Electrode
Consists of a platinum electrode immersed in a
solution with a hydrogen-ion concentration of
1M Solution is at 25 C and the electrode is a
small square of platinum foil coated with finely
divided platinum, known as platinum
black. Hydrogen gas at 101 kPa is bubbled around
the platinum electrode.
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Standard Hydrogen Electrode
2H (aq, 1M) 2e- H2 (g, 101kPa) EOH
0.00V Double arrows indicate the reaction is
reversible. Standard reduction potential of H is
the tendency of H ions to acquire electrons and
be reduced to H2 (g) Whether the half-cell
reaction occurs as a reduction or as an oxidation
is determined by the reduction potential of the
half-cell to which the standard hydrogen
electrode is connected.
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Standard Reduction Potentials
A voltaic cell can be made by connecting a
standard hydrogen half-cell to a standard zinc
half-cell.
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Standard Reduction Potentials
To determine the overall reaction for this cell,
first identify the half-cell in which reduction
takes place. In all electrochemical cells,
reduction takes place at the cathode and
oxidation takes place at the anode. A voltmeter
gibes a reading of 0.76 V when the zinc
electrode is connect to the negative terminal and
the hydrogen electrode is connect to the positive
terminal. Zinc is oxidized anode and Hydrogen
ions are reduced hydrogen electrode is the
cathode. .
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Standard Reduction Potentials
Oxidation Zn (s) Zn2 (aq) 2e- (at
anode) Reduction 2H(aq) 2e- H2 (g)
(at cathode) You can determine the standard
reduction potential of a half-cell by using a
standard hydrogen electrode and the equation for
standard cell potential E0cell E0red -
E0oxid E0cell E0H - E0Zn2 0.76 V
0.00 V - E0Zn2 E0Zn2 -0.76 V
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Standard Reduction Potentials
The standard reduction potential for the zinc
half-cell is -0.76 V. The value is negative
because the tendency of zinc ions to be reduced
to zinc metal in this cell is less than the
tendency of hydrogen ions to be reduced to
hydrogen gas. Consequently, the zinc ions are
not reduced. Instead, the opposite occurs Zinc
metal is oxidized to zinc ions.
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Standard Reduction Potentials
For a standard copper half-cell, the measured
standard cell potential is 0.34V. Copper is the
cathode and Cu2 ions are reduced to Cu
metal Hydrogen half-cell is the anode, and H2 gas
is oxidized to H ions. E0cell E0red -
E0oxid E0cell E0Cu2 - E0H 0.34 V
E0Cu2 - 0.00 E0Cu2 0.34 V
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Discussion
The two half-cells of a voltaic cell are
competing for electrons. Oxidation or reduction
could occur in either cell. The half-cell with
the more positive reduction potential will win
the competition and undergo reduction. The
potential produced by the electrochemical cell is
the difference in the reduction potentials of the
two half-cell reactions. The quantitative value
of any half-cell potential is obtained by
measuring it against the standard hydrogen
electrode. (Review table 21.2 page 674)
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Discussion
Which reactions have the greatest tendency to
occur as reductions? Activity series of metals
have the most active metals at the top. Because
active metals lose electrons easily, they are
most easily oxidized. Thus, the ions of active
metals are least likely to be reduced. The
potential produced by the electrochemical cell is
the difference in the reduction potentials of the
two half-cell reactions.
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Calculating Standard Cell Potentials
To function, a cell must be constructed of two
half-cells. The half-cell reaction having the
more positive (or less negative) reduction
potential occurs as a reduction in the cell. You
can use the know standard reduction potentials
for various half-cells to predict the half-cell
in which reduction and oxidation will occur. If
the cell potential for a given redox reaction is
then the reaction is spontaneous as written. If
the cell potential is - , then the reaction in
nonspontaneous.
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Question
Determine whether the following redox reaction
will occur spontaneously. 3Zn2(aq) 2Cr(s)
3Zn(s) 2Cr3(aq) Oxidation Cr(s)
Cr3(aq) 3e- E0Cr3 -0.74V Reduction
Zn2(aq) 2e- Zn(s) E0Zn2
-0.76V E0cell E0red - E0oxid E0cell
E0Zn2 - E0Cr3 E0cell -0.76
(-0.74) E0cell -0.02 V (nonspontaneous)
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Question
Is this redox reaction spontaneous as
written? Co2(aq) Fe(s) Co(s)
Fe2(aq) Oxidation Fe(s) Fe2(aq)
2e- E0Cr3 -0.44V Reduction Co2(aq)
2e- Co(s) E0Zn2 -0.28V E0cell
E0red - E0oxid E0cell E0Co2 -
E0Fe2 E0cell -0.28 (-0.44) E0cell 0.16 V
(spontaneous)
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Question
Determine the cell reaction for a voltaic cell
composed of the following half-cells. Fe3(aq)
e- Fe2(aq) E0Fe3 0.77V Ni2(aq)
2e- Ni(s) E0Ni2 -0.25V The
half-cell with the more positive reduction
potential is the one in which reduction occurs
(the cathode) Oxidation Ni(s) Ni2(aq)
2e- Reduction 2Fe3(aq) 2e-
2Fe2(aq) (balance e-) Ni(s)
2Fe3(aq) Ni2(aq) 2Fe2(aq)
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Discussion
In the previous example, we had to multiply the
Fe half-cell reaction by a factor of 2 to cancel
out the electrons. Even though, there were two
times as many electrons present, the tendency for
the electrons to flow is not two times greater.
The tendency, which is measured by the E0 value,
remains the same.
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Questions
A voltaic cell is constructed using the following
half reactions. Cu2(aq) 2e- Cu(s)
E0Cu2 0.34V Al3(aq) 3e- Al(s)
E0Al3 -1.66V 2Al(s) 3Cu2(aq)
2Al3(aq) 3Cu(s) A voltaic cell is
constructed using the following half
reactions. Ag(aq) e- Ag(s)
E0Ag 0.80V Cu2(aq) 2e- Cu(s)
E0Cu2 0.34V Cu(s) 2Ag(aq)
Cu2(aq) 2Ag(s)
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Question
Calculate the standard cell potential for the
Ni/Fe voltaic cell. Half-reactions are as
follows Fe3(aq) e- Fe2(aq) E0Fe3
0.77V Ni2(aq) 2e- Ni(s)
E0Ni2 -0.25V E0cell E0Fe3 -
E0Ni2 E0cell 0.77 V (-0.25 V) E0cell
1.02 V
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Question
A voltaic cell is constructed using the following
half-reactions Al3(aq) 3e- Al(s)
E0Al3 -1.66V Cu2(aq) 2e- Cu(s)
E0Cu2 0.34V E0cell E0Cu2 -
E0Al3 E0cell 0.34 V (-1.66 V) E0cell
2.00 V
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Question
A voltaic cell is constructed using the following
half-reactions Ag(aq) e- Ag(s)
E0Ag 0.80V Cu2(aq) 2e- Cu(s)
E0Cu2 0.34V E0cell E0Ag -
E0Cu2 E0cell 0.80 V (0.34 V) E0cell
0.46 V
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Questions
What causes the electrical potential of a
cell? Competition for electrons between two
half-cells What is the electrical potential of a
standard hydrogen electrode? Assigned a value of
0.00 V at 25ºC How can you find the standard
reduction potential of a half-cell? By
connecting it to a standard hydrogen electrode
and measuring the cell potential What cell
potential values indicate a spontaneous reaction?
A nonspontaneous reaction? Positive cell
potential - spontaneous
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Homework
Using the reduction potentials from table 21.2,
create an electrochemical cell that will operate
spontaneously. Calculate the cell potential
Write the equations for the two half-reactions
and the overall cell reaction. Use the shorthand
method to represent the cell
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End of section 20.2
71
Electrolytic Cells
An electric current can be used to make a
nonspontaneous redox reaction go forward.
Electrolysis the process in which electrical
energy is used to bring about a nonspontaneous
chemical change. Examples of electrolysis are
silver-plated dishes and utensils, gold-plated
jewelry, and chrome-plated automobile parts.
Electrolytic cell the apparatus in which
electrolysis is carried out is an electrochemical
cell used to cause a chemical change through the
application of electrical energy.
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Electrolytic Cells
An electrolytic cell uses electrical energy
(direct current) to make a nonspontaneous redox
reaction proceed to completion. In both voltaic
and electrolytic cells, electrons flow from the
anode to the cathode in the external circuit. For
both types of cells, the electrode at which
reduction occurs is the cathode. The key
difference between voltaic and electrolytic cells
is that in a voltaic cell, the flow of electrons
is the result of a spontaneous redox reaction,
whereas in an electrolytic cell, electrons are
pushed by an outside power source, such as a
battery.
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Voltaic cell energy is released from a
spontaneous redox reaction. Electrolytic cell -
energy is absorbed to drive a non-spontaneous
reaction.
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Electrolytic Cells
Electrolytic and voltaic cells also differ in the
assignment of charge to the electrodes. In an
electrolytic cell, the cathode is considered to
be the negative electrode, because it is
connected to the negative electrode of the
battery. The anode in an electrolytic cell is
considered to be the positive electrode because
it is connected to the positive electrode of the
battery. In a voltaic cell, the anode is the
negative electrode and the cathode is the
positive electrode.
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Electrolytic Cells
Electrolytic processes are used to separate
active metals such as aluminum, magnesium, and
sodium from their salts. The same process is used
to recover metals from ores.
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Electrolysis of Water
When a current is applied to two electrodes
immersed in pure water, nothing happens. When an
electrolyte such as sulfuric acid or potassium
nitrate in low concentration is added to the pure
water, the solution conducts electricity and
electrolysis occurs. The products of the
electrolysis of water are hydrogen gas and oxygen
gas.
77
Electrolysis of Water
Water is reduced to hydrogen at the
cathode Reduction 2H2O(l) 2e- H2(g)
2OH-(aq) Water is oxidized at the
anode Oxidation 2H2O(l) O2(g) 4H (aq)
4e- The region around the anode turns acidic
due to an increase in H ions. The region around
the cathode turns basic due to the production of
OH- ions.
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Electrolysis of Water
The overall cell reaction 4H2O(l) 4e-
2H2(g) 4OH-(aq) (x2 to balance) 2H2O(l)
O2(g) 4H (aq) 4e- 6H2O(l) 2H2(g)
4OH-(aq) O2(g) 4H(aq) The ions produced
tend to recombine to form water, so they are not
included in the net reaction. 6H2O(l)
electrolysis H2(g) O2(g)
79
Electrolysis of Brine
If the electrolyte in an aqueous solution is more
easily oxidized or reduced than water, then the
products of electrolysis will be substances other
than hydrogen and oxygen. Example is brine (a
concentrated aqueous solution of sodium chloride)
which produces chlorine gas, hydrogen gas, and
sodium hydroxide. During electrolysis of brine,
chloride ions are oxidized to produce chlorine
gas at the anode. Oxidation 2Cl-(aq)
Cl2(g) 2e- (at anode)
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Electrolysis of Brine
Water is reduced to produce hydrogen gas at the
cathode. Reduction 2H2O(l) 2e- H2(g)
2OH-(aq) (at cathode) Sodium ions are not reduced
to sodium metal in the process because water
molecules are more easily reduced than are sodium
ions. The reduction of water produces hydroxide
ions as well as hydrogen gas. Thus the
electrolyte in solution becomes sodium hydroxide.
81
Electrolysis of Brine
The overall ionic equation 2H2O(l) 2e-
H2(g) 2OH-(aq) 2Cl-(aq) Cl2(g)
2e- 2H2O(l) 2Cl-(aq) H2(g) 2OH-(aq)
Cl2(g) The spectator ion Na can be included in
the equation (as part of NaCl and of NaOH) to
show the formation of sodium hydroxide during the
electrolytic process 2NaCl (aq) 2H2O(l)
2Cl-(aq) H2(g) 2NaOH(aq) Cl2(g)
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Electrolysis in Metal Processing
Electrolytic cells are commonly used in the
plating, purifying and refining of metals. Many
of the shiny, metallic objects you see every day,
such as chrome-plated fixtures or nickel-plated
coins, were manufactured with the help of
electrolytic processes. Electroplating is the
deposition of a think layer of metal on an object
in an electrolytic cell. An object may be
electroplated to protect the surface of the base
metal from corrosion or to make it more
attractive.
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Electrolysis in Metal Processing
An object that is to be silver-plated is made the
cathode in an electrolytic cell. The anode is
the metallic silver that is to be deposited The
electrolyte is a solution of a silver salt, such
as silver cyanide. When a direct current is
applied, silver ions move from the anode to the
object to be plated. Reduction Ag (aq)
e- Ag (s) (at cathode) The net result
is that silver transfers from the silver
electrode to the object being plated.
84
Electrolysis in Metal Processing
Many factors contribute to the quality of the
metal coating that forms. In the plating
solution, the concentration of the cations to be
reduced must be carefully controlled. The
solution must also contain compounds to control
the acidity and to increase the conductivity.
Other compounds may be used to make the metal
coating brighter or smoother.
85
Electrolysis in Metal Processing
Electroforming is a process in which an object
is reproduced by making a metal mold of it at the
cathode of a cell. A phonograph record can be
coated with metal so it will conduct a current.
It is then electroplated with a thick coating of
metal. This coating can be stripped off and used
as a mold to produce copies of the record.
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Electrowinning
Electrowinning a process where impure metals
can be purified in electrolytic cells. The
cations of molten salts or aqueous solutions are
reduced at the cathode to give very pure metals.
A common use is in the extraction of aluminum
form its ore, bauxite. (Al2O3) In a method know
as the Hall-Heroult process, purified alumina is
dissolved in molten cryolite (Na3AlF6), and
heated to above 1000ºC in a carbon line tank.
87
Electrowinning
The carbon lining, connected to a direct current,
serves as the cathode. The anode consists of
carbon rods dipped into the tank. At the
cathode, Al3 ions are reduced, forming molten
aluminum. At the anode, carbon is oxidized,
forming carbon dioxide gas. 2Al2O3(l) 3C(s)
4Al(l) 3CO2(g)
88
Other Electrolytic Processes
Electrorefining a piece of impure metal is made
the anode of the cell. It is oxidized to the
cation and then reduced to the pure metal at the
cathode. Electrorefining technique is used to
obtain ultrapure silver, lead and copper. Other
electrolytic processes are centered on the anode
rather than the cathode. Electropolishing - the
surface of an object at the anode is dissolved
selectively to give it a high polish.
Electromachining - a piece of metal at the anode
is partially dissolved until the remaining
portion is an exact copy of the object at the
cathode.
89
Questions
What is the difference between an electrolytic
cell and a voltaic cell? Voltaic cell uses an
electrochemical reaction to produce electrical
energy. An electrolytic cell uses electrical
energy to bring about a chemical change. What
products form during the electrolysis of water?
H2 (g) and O2 (g) What chemical changes occur
during the electrolysis of brine? Chloride ions
are oxidized to produce chlorine gas and water is
reduced to produce hydrogen gas.
90
Questions
What are some application of electrolysis in the
field of metallurgy? Electroplating (deposition
of a thin layer of metal on an object),
electrorefining (purification of metals ) and
electrowinning (extraction of metals) What is the
charge on the anode of an electrolytic cell? Of a
voltaic cell? Electrolytic cell anode ()
voltaic cell anode (-) Which process, oxidation
or reduction, always occurs at the cathode of an
electrolytic cell? Reduction
91
Questions
Can metallic sodium be obtained by electrolyzing
brine? No the products are chlorine gas,
hydrogen gas, and sodium hydroxide. Sodium is
obtained by electrolysis of molten NaCl in the
Downs cell, which operates at 801ºC to keep it
melted. The anode and cathode of the cell are
separated to prevent recombination of sodium and
chlorine. Reduction of Na occurs at a graphite
anode. Liquid sodium rises to the top of the
molten NaCl and is drawn off. 2NaCl (l)
2Na (l) Cl2 (g)
92
End of Chapter 20