Electrochemistry
Electrochemical reactions are oxidation-reduction reactions.
| The two parts of the reaction are physically separated.
Chapter 21a
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Electrochemistry: The
Electrolytic Cell
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The oxidation reaction occurs in one cell.
The reduction reaction occurs in the other cell.
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Electrochemistry
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1.
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There are two kinds electrochemical cells.
Electrochemical cells containing spontaneous
chemical reactions are called voltaic or galvanic
cells.
„
2.
Electrical Conduction
Metals conduct electric currents well in a process called
metallic conduction.
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The generation of electric current from a chemical
reaction.
In metallic conduction there is electron flow with no atomic
motion.
Metal atoms changing oxidation states without moving.
• E.g. Oxidative phosphorylation
Electrochemical cells containing in nonspontaneous
chemical reactions are called electrolytic cells.
„
The use of electric current to produce a chemical change.
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Electrical Conduction
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Electrodes
The following convention for electrodes is correct for
either electrolytic or voltaic cells:
| The cathode is the electrode at which reduction
occurs.
In ionic or electrolytic conduction ionic motion
transports the electrons.
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Positively charged ions, cations, move toward the
negative electrode.
Negatively charged ions, anions, move toward the
positive electrode.
• The cathode is negative in electrolytic cells and positive
in voltaic cells.
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The anode is the electrode at which oxidation
occurs.
• The anode is positive in electrolytic cells and negative in
voltaic cells.
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Electrodes
Electrolytic Cells
Inert electrodes do not react with the liquids or
products of the electrochemical reaction.
Two examples of common inert electrodes are
graphite and platinum.
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Electrical energy is used to force nonspontaneous
chemical reactions to occur.
| The process is called electrolysis.
| Two examples of commercial electrolytic reactions
are:
1.
2.
The electroplating of jewelry and auto parts.
The electrolysis of chemical compounds.
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Electrolytic Cells
Figure 11.19:
Voltaic Cell
Electrolytic Cell
Electrolytic cells consist of:
|
1.
2.
3.
A container for the reaction mixture.
Two electrodes immersed in the reaction mixture.
A source of direct current.
Electrolytic cells uses electrical energy to produce a
chemical change.
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The electrical energy forces a current through a cell that has a
negative potential.
The electrical energy forces a chemical change to occur.
(a) A standard galvanic cell
(b) A standard electrolytic cell
The cell in (b) has a power source that forces the electrons in the
opposite direction from the voltaic cell in (a).
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Counting Electrons: Coulometry and Faraday’s
Law of Electrolysis
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Counting Electrons: Coulometry and Faraday’s Law of
Electrolysis
The stoichiometry of electrolysis processes can
quantify “how much chemical change occurs with the flow of
a given current for a specific time”.
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Faraday’s Law - The amount of substance
undergoing chemical reaction at each electrode
during electrolysis is directly proportional to the
amount of electricity that passes through the
electrolytic cell.
A faraday is the amount of electricity that reduces
one equivalent of a species at the cathode and
oxidizes one equivalent of a species at the anode.
1 faraday of electricity ≡ 6.022 × 1023 e-
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Counting Electrons: Coulometry and Faraday’s Law of
Electrolysis
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Counting Electrons: Coulometry and Faraday’s Law of
Electrolysis
A coulomb is the amount of charge that passes a
given point when a current of one ampere (A) flows
for one second.
1 ampere (amp) = 1 coulomb/second
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Faraday’s Law states that during electrolysis, one
faraday of electricity (96,485 coulombs) reduces and
oxidizes, respectively, one equivalent of the
oxidizing agent and the reducing agent.
This corresponds to the passage of one mole of electrons
through the electrolytic cell.
1 faraday ≡ 6.022 × 1023 e-
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1 faraday ≡ 1.0 mol e –
1 equivalent of oxidizing agent ≡ gain of 6.022 ×1023 e –
1.0 mol e ≡ 96, 485 coulombs
1 equivalent of reducing agent ≡ loss of 6.022 ×1023 e –
–
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Counting Electrons: Coulometry and Faraday’s Law of
Electrolysis
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The Electrolysis of Water
Example: Calculate the mass of palladium produced by the
reduction of palladium (II) ions during the passage of 3.20
amperes of current through a solution of palladium (II)
sulfate for 30.0 minutes.
Cathode :
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Pd 2+ + 2e- → Pd 0
1 mol 2 mol 1 mol
106 g 2(96, 485) 106 g
C
s
1 mol e− = 96, 485 C
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3.20 amp = 3.20
?g =
( 30.0
Hydrogen and oxygen combine spontaneously to
form water.
|
−
 60 s   3.20 C   1 mol e   mol Pd  106 g Pd
min ) 
= 3.16 g Pd



- 
 min   s   96, 485 C   2 mol e  mol Pd
The decrease in free energy that accompanies this
spontaneous reaction can be used to run fuel cells to
produce electricity.
The reverse process, which is not spontaneous,
requires energy to occur.
The formation of oxygen and hydrogen gases from
water can be forced by electrolysis.
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Counting Electrons: Coulometry and Faraday’s Law of
Electrolysis
The Electrolysis of Water
Anode reaction
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2 H 2O → O2( g ) + 4 H + + 4 e –
Cathode reaction 2(2 H 2O + 2e – → H 2( g ) + 2OH – )
Cell reaction 6 H 2O → 2 H 2( g ) + O2( g ) + 4
H + +
4 OH
–
Example: Calculate the volume of oxygen (measured at STP)
produced by the oxidation of water during the passage of 3.20
amperes of current for 30.0 minutes.
Anode : 2 H 2O → O2( g ) + 4 H + + 4e 2 mol
4 H 2O
1 mol
4 mol 4 mol
22.4 LSTP
The overall reaction is 2 H 2O → 2 H 2( g ) + O2( g )
? LSTP O2 = 3.20 amp = 3.20
4 ( 96,500 C )
C
s
−
1 mol e = 96, 485 C
1.0 mol = 22.4 LSTP
( 30.0
−
 60 s   3.20 C   1 mol e   mol O2   22.4 LSTP O2 
min ) 
=



- 
 min   s   96, 485 C   4 mol e   mol O2 
= 0.334 LSTP O2 or 334 mLSTP O2
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The Electrolysis of Molten Sodium Chloride
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Liquid sodium is produced at one electrode.
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The Electrolysis of Molten Sodium Chloride
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Indicates that the reaction Na+(A) + e- → Na(A) occurs at
this electrode.
Is this electrode the anode or cathode?
Reduction occurs at the cathode.
In all electrolytic cells, electrons are forced to flow
from the positive electrode (anode) to the negative
electrode (cathode).
Gaseous chlorine is produced at the other electrode.
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Indicates that the reaction 2 Cl- → Cl2(g) + 2 e- occurs at
this electrode.
Is this electrode the anode or cathode?
Oxidation occurs at the anode.
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The Electrolysis of Molten Sodium Chloride
The Electrolysis of Molten Sodium Chloride
Diagram of this electrolytic cell.
ee–
- electrode
e–
gains
generating
liquid Na0.
Na (A)
e–
Na+ +
Na+ + e- → Na(A)
cathode reaction
Cl2
2e–
2Cl–
chloride loses
generating Cl2
gas
The nonspontaneous redox reaction that occurs is:
Anode reaction
(
+ electrode
e–
e–
Na+
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Generator-source
of DC
e-
2 Cl - → Cl2( g ) + 2 e-
Cathode reaction 2 Na + + e- → Na( A )
e–
Cell reaction 2 Cl + 2 Na
-
+
)
→ Cl2( g ) + 2 Na( A )
molten NaCl
Porous barrier
2Cl- → Cl2 (g) + 2eanode reaction
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Figure 11.25: The Downs cell for the electrolysis of molten sodium
chloride.
Sodium metal is
produced by the
electrolysis of molten
sodium chloride.
NaCl is mixed with
CaCl2 to lower the
melting point (from
800oC to 600oC).
The liquid sodium is
drained, cast into
blocks and stored in
inert solvents.
The Electrolysis of Aqueous Sodium Chloride
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In this electrolytic cell, hydrogen gas is produced at
one electrode.
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The aqueous solution becomes basic near this electrode.
What reaction is occurring at this electrode? Gaseous
chlorine is produced at the other electrode.
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The Electrolysis of Aqueous Potassium Chloride
The Electrolysis of Aqueous Potassium Chloride
Cell diagram
– pole of battery
What reaction is occurring at this electrode?
e- flow
+ pole of battery
Battery, a source
of direct current
e- flow
These experimental facts lead us to the following
nonspontaneous electrode reactions:
Anode reaction
- electrode
2 Cl – → Cl2( g ) + 2 e –
Cathode reaction 2 H 2O + 2 e – → H 2( g ) + 2 OH –
+ electrode
H2 gas
Cl2 gas
Cell reaction 2 Cl – + 2 H 2O → H 2( g ) + Cl2( g ) + 2 OH –
Na + is a spectator ion. Note that water is electrolyzed !
aqueous NaCl
2 H2O + 2e- → H2 (g) + 2 OHcathode reaction
2Cl- → Cl2 (g) + 2eanode reaction
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Electrolytic Cells
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Commercial Applications of Electrolytic Cells
Electrolytic Refining and Electroplating of Metals
In all electrolytic cells the most easily reduced
species is reduced and the most easily oxidized
species is oxidized.
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Impure metallic copper can be purified
electrolytically to ≈ 100% pure Cu.
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The impurities commonly include some active metals
plus less active metals such as: Ag, Au, and Pt.
The cathode is a thin sheet of copper metal
connected to the negative terminal of a direct
current source.
The anode is large impure bars of copper.
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Commercial Applications of Electrolytic Cells
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Commercial Applications of Electrolytic Cells
The electrolytic solution is CuSO4 and H2SO4
The impure Cu dissolves to form Cu2+.
The Cu2+ ions are reduced to Cu at the cathode.
Anode ( impure
)
Cathode ( very pure
Net rxn.
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This effectively removes them from the Cu metal.
Zn 0 → Zn 2 + + 2e −
Cu (0s ) → Cu (2aq+ ) + 2e −
)
Any active metal impurities are oxidized to cations
that are more difficult to reduce than Cu2+.
Fe0 → Fe2 + + 2e −
And so forth for other
Cu (2+aq ) + 2e− → Cu (0s)
No net rxn.
active metals
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Figure 11.24: Schematic of the electroplating of a
spoon.
Commercial Applications of Electrolytic Cells
The spoon is the cathode and
is plated out by the Ag+ ions
that are released from the solid
silver bar that is the anode.
Metal Plating
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Objects can be plated by making a particular object
a cathode in a tank with ions of the plating metal.
A salt bridge is not required
because Ag+ are at acting at
both electrodes.
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Copper Plating
Commercial Applications of Electrolytic Cells
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The less active metals are not oxidized and
precipitate to the bottom of the cell.
These metal impurities can be isolated and separated
after the cell is disconnected.
Some common metals that precipitate include:
Ag , Au, Pt , Pd
( Se, Te )
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Corrosion
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Corrosion
Metallic corrosion is the oxidation-reduction reactions
of a metal with atmospheric components such as CO2,
O2, and H2O.
Metals corrode because they oxidize easily.
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Corrosion of iron
The importance of steel in many of our structures,
controlling corrosion is a very important issue.
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Many common metals that are used for structural and
decorative purposed have standard reduction potentials that
are more negative and oxygen.
The corrosion mechanism involves electrochemical
processes.
4 Fe0 + 3 O20 → 2 Fe2O3 ( overall reaction )
The reaction occurs rapidly at exposed points.
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Corrosion
steel
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The surface of steel is not uniform.
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The chemical composition of steel is not a homogeneous
mixture.
Stress points are produced on the surface due to physical
strains.
At these stress areas, iron can be more easily oxidized in
some regions than in other regions.
o Oxidized areas act as anodes
o The other areas act as cathodes
Fe2+ ions travel through the surface moisture to the region
acting as a cathode.
In the region of the cathode, the Fe2+ ions react with O2 to
form rust.
The moisture acts as a salt bridge in the process of corrosion.
Without moisture, steel does not rust.
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Corrosion Protection
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1.
Corrosion Protection
Some examples of corrosion protection.
Plate a metal with a thin layer of a less active (less easily
oxidized) metal.
2. Galvanizing, the coating of steel with zinc, provides
a more active metal on the exterior.
The thin coat of Zn must be oxidized before Fe begins to rust.
"Tin plate " or " chromium plate " for steel.
Zinc
Steel
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Corrosion Protection
Corrosion Protection
3. Connect the metal
to a sacrificial
anode, a piece of a
more active metal.
Soil pipes and ship hulls have Mg and Zn on the exterior as sacrificial anodes.
Magnesium is easily
oxidized; protecting the iron
from oxidation.
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Corrosion Protection
Corrosion Protection
4. Allow a protective film to form naturally.
5.
4 Al 0 + 3 O20 → 2 Al2O3
Paint or coat with a polymeric material such as
plastic or ceramic.
Steel bathtubs are coated with ceramic.
Al2O3 forms a hard , transparent film on exterior of aluminum foil.
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End of Chapter 11b
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Electrochemistry is an important part of the
electronics industry.
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