Chapter 18 Electrochemistry
The batteries we use to power
portable computers and other
electronic devices all rely on
redox reactions to generate
an electric current. Redox
reactions are central to the
development of small, light,
long-lasting power sources.
To a large extent, the future
development of technology
depends on the capabilities
of these power sources. In
this chapter, we see what is
involved in using chemical
reactions to generate
electricity.
Assignment for Chapter 18
18.2; 18.6; 18.17; 18.27
Electrochemistry in Biological Systems
Ion transport
Muscle contraction
Neuron excitation
Psychological activities
Electrochemistry
• Electrochemistry is a branch of chemistry
that deals with the use of spontaneous
chemical reactions to produce electricity
and the use of electricity to bring about
spontaneous chemical change.
Why electricity and chemical reaction may be connected?
Because both involve electrons! The motion or transfer of
electrons!
Proton transfer: acid-base reactions
Electron transfer: redox reactions
Half-Reactions
• Chemical equations showing the changes involved
only in oxidation or reduction.
Oxidation half-reaction:
2
Mg(s)  Mg (aq)  2e
Reduction half-reaction:
Fe (aq)+3e  Fe(s)
2+
-

Example
The oxidation of iron(II) to iron(III):
Fe (aq)  Fe (aq)+e
2+
3+
-
The reduction of copper(II) to copper metal:
Cu (aq)+2e  Cu(s)
2+
-
Exercise:
Aluminum metal is oxidized to Al3+ in aqueous solution:
Al(s)  Al (aq)+3e
3+
-
Balancing Redox Reactions
MnO4-(aq) + H2C2O4 (aq)  Mn2+ (aq) + CO2 (g)
!
• The chemical equation of a reduction halfreaction is added to that of an oxidation
half-reaction to form the chemical equation
for the overall redox reaction.
Use H2O and H+ and OH- as “extra” chemicals
Balancing Redox Reactions
in acidic solutions
MnO4-(aq) + H2C2O4 (aq)  Mn2+ (aq) + CO2 (g)
!
Reduction: MnO4- Mn2+ (aq)
Oxidation: H2C2O4 (aq)  CO2 (g)
Balance all elements except H and O:
Reduction: MnO4- Mn2+ (aq)
Oxidation: H2C2O4 (aq)  2CO2 (g)
Balance O using water and H using H+:
Reduction: MnO4- + 8H+ Mn2+ (aq)+4H2O
Oxidation: H2C2O4 (aq)  2CO2 (g) + 2H+
Balancing Redox Reactions
Balance Charge:
Reduction: MnO4- + 8H+ +5e- Mn2+ (aq)+4H2O
Oxidation: H2C2O4 (aq)  2CO2 (g) + 2H+ + 2eReduction x 2 + Oxidation x 5
2MnO4- + 16H+ +10e- 2Mn2+ (aq)+8H2O
5H2C2O4 (aq)  10CO2 (g) + 10H+ + 10e2MnO4- (aq)+ 5H2C2O4 (aq) +6H+ (aq) 
2Mn2+ (aq)+8H2O(l) + 10CO2 (g)
Figure 18.1 A schematic diagram of how to balance
a redox equation by balancing the half-reactions
separately and then combining them.
Example
• Write down the balanced net ionic reaction for the
reaction: Cu+HNO3Cu2++NO.
Cu+H++NO3-Cu2++NO
CuCu2+
2NO3-2NO + 4H2O
CuCu2+
8H++2NO3-2NO + 4H2O
CuCu2++2e-
8H++2NO3-+6e-2NO + 4H2O
x3
3Cu(s)+2NO3-(aq)+8H+(aq)3Cu2+(aq)+2NO(g)+4H2O(l)
Exercise
MnO4++H2SO3  HSO4- + Mn2+.
Write down the balanced net ionic reaction.
Balancing Redox Reactions
in basic solutions
Br- (aq)+ MnO4-(aq) MnO2(s)+BrO3-(aq)
Br- BrO3-
MnO4-(aq) MnO2(s)
Br- + 3H2OBrO3-
MnO4-(aq) MnO2(s) +2H2O
Br- + 3H2O+6OH-BrO3-+6H2O
MnO4-(aq)+4H2O MnO2(s) +2H2O+4OHBr- + 6OH-BrO3-+3H2O +3eMnO4-(aq)+2H2O MnO2(s) +4OH-+6e2MnO4-(aq)+ Br-(aq) +H2O(l) MnO2(s)+BrO3-(aq)+2OH-(aq)
Electrochemical Cell
• A device in which an electric current is
either produced by a spontaneous chemical
reaction or is used to bring about a
nonspontaneous reaction.
A galvanic cell is an electrochemical cell in which a spontaneous
chemical reaction is used to generate an electric current.
Figure 18.2 In an
electrochemical cell, a reaction
takes place in two separate
negative
regions. Oxidation occurs at
one electrode, and the
electrons released travel
through the external circuit to
the other electrode, where they
cause reduction. The site of
oxidation is called the anode,
and the site of reduction is
called the cathode.
positive
Any two objects that have different
(first) ionization energies may function
as a cell.
- 1.234 +
- 0.02 +
Figure 18.3(a) When a bar of zinc is placed in a beaker of copper(II)
sulfate solution, copper is deposited on the zinc and the blue copper
(II) ions are gradually replaced by colorless zinc ions. (b) The residue
in the beaker is copper metal. No more copper ions can be seen in
solution.
Zn(s)+Cu (aq)  Zn (aq)+Cu(s)
2+
2+
Figure 18.4 The reaction shown in Fig. 18.3 takes
place all over the surface of the zinc as electrons
are transferred to the Cu2 ions in solution.
Cu 2+ (aq)+2e-  Cu(s)
Zn(s)  Zn (aq)+2e
2+
-
Figure 18.5 The Daniell cell consists of copper and zinc electrodes
dipping into solutions of copper(II) sulfate and zinc sulfate, respectively.
The two solutions make contact through the porous pot, which allows
ions to pass through to complete the electrical circuit.
Zn(s)|Zn 2+ Cu 2+ |Cu(s)
Zn(s)  Zn 2+ (aq)+2e-
Cu (aq)+2e  Cu(s)
2+
-
Electrodes and Cell Diagram
Zn(s)|Zn 2+ Cu 2+ |Cu(s)
+
H (aq)|H 2 (g)|Pt(s)
3+
2+
Fe (aq),Fe (aq) |Pt(s)
2+
2+
Zn(s)|Zn (aq)|Cu (aq) |Cu(s)
Figure 18.6 This cell is typical of galvanic cells used in the
laboratory. The two electrodes are connected by an external circuit
and a salt bridge. The latter completes the electrical circuit within the
cell.
Pt(s)|Fe3+ (aq),Fe2+ (aq) ||Cu 2+ (aq) |Cu(s)
Figure 18.7 The cell potential is measured with an electronic voltmeter, a
device that draws negligible current so that the composition of the cell
does not change during the measurement. The display shows a positive
value when the  terminal of the meter is connected to the cathode of the
galvanic cell.
2+
2+
Zn(s)|Zn (aq)||Cu (aq) |Cu(s)
Cell Potential
E = 1.1 V
The cell potential
• An indication of the electron-pulling and –
pushing power of the cell reactions; cell
reactions at equilibrium generate zero
potential.
Figure 18.8 Electrons produced by oxidation leave a galvanic cell at the
anode (), travel through the external circuit, and reenter the cell at the
cathode (), where they cause reduction. The circuit is completed inside
the cell by migration of ions through the salt bridge. A salt bridge is
unnecessary when the two electrodes share a common electrolyte.
positive
negative
Figure 18.9 This schematic picture of a galvanic cell indicates the
identities of the anode and cathode, displays the oxidation and reduction
half-reactions, and shows the direction of electron flow.
Describing a galvanic cell and identifying the cell reaction
Hg 2+ (aq)+2e-  2Hg(l) 2Hg(l)+2Cl- (aq)  Hg 2Cl2 (s)+2e-
 Hg (aq)+2Cl (aq)  Hg 2Cl2 (s)
2+
-
Hg(l)|Hg 2Cl 2 (s)|HCl(aq)||Hg (NO ) (aq) |Hg(l)
2
3 2
(KCl gel)
Exercise: Describing a galvanic cell and identifying the cell reaction
(Assume platinum electrode is used)
H 2 (g)  2H + (aq)+2e-
Co3+ (aq)+e-  Co2+ (aq)
H2 (g)+2Co (aq)  2H (aq)+ 2Co (aq)
3+
+
+
3+
2+
2+
Pt(s)|H 2 (g)||H (aq)||Co (aq),Co (aq) |Pt(s)
Figure 18.12 The cell potential can be thought of as being the difference
of the two reduction potentials produced by the two electrodes. The cell
potential is positive if the cathode has a higher potential than the anode.
Cell potential and electrode potential
standard cell potential:
E 0  E 0 (cathode)  E 0 (anode)
Fe(s)|Fe 2+ (aq)||Ag + (aq) | Ag(s):
E 0  1.24 V
Hydrogen standard potential
2H +2e  H 2 (g): E  0 V
+
-
0
Zn(s)|Zn (aq)||H (aq) | H 2 (g)|Pt(s): E  0.76 V
2+
+
0
E 0  E 0 (H + / H 2 )  E 0 (Zn 2+ / Zn)  0.76 V
 E 0 (Zn 2+ / Zn)  0.76 V
Figure 18.13 The variation of standard potentials in the main groups of
the periodic table. Note that the most negative values occur in the s block
and the most positive values occur close to fluorine.
Example: deducing the standard potential of an electrode
The standard potential of a Zn2+/Zn electrode is -0.76 V and
the standard potential of the cell
Zn(s)|Zn2+(aq)||Cu2+(aq)|Cu(s) is 1.10 V.
What is the standard potential of Cu2+/Cu electrode?
E 0 (Cu 2+ /Cu)  E 0 (Zn 2+ /Zn)  1.10 V
E (Cu / Cu )  1.10 V  E (Zn / Zn)
 1.10 V - 0.76 V=+0.34 V
0
2+
0
2+
Exercise: deducing the standard potential of an electrode
The standard potential of a Fe2+/Fe electrode is -0.44 V and
the standard potential of the cell
Fe(s)|Fe2+(aq)||Pb2+(aq)|Pb(s) is 0.31 V.
What is the standard potential of Pb2+/Pb electrode?
E 0 (Pb2+ /Pb)  E 0 (Fe2+ /Fe)  0.31 V
E ( Pb / Pb)  0.31 V+E ( Fe / Fe)
 0.31 V  0.44 V =-0.13 V
0
2+
0
2+
Figure 18.14 The significance of standard potentials. Only couples with
negative standard potentials (and hence lying below hydrogen) can
reduce hydrogen ions to hydrogen gas. The reducing power increases as
the standard potential becomes more negative.
E 0 (Cu 2+ / Cu)=+0.34 V
E (Cu /Cu)  E (H /H)  0.34 V
0
2+
0
+
Cu (aq)+H 2 (g)  Cu(s)+H (aq)
2+
+
E 0 (Zn2+ / Zn)=-0.76 V
E (Zn /Zn)  E (H /H)  0.76 V
0
2+
0
+
Zn(s)+H (aq)  Zn (aq)+H 2 (g)
+
2+
Figure 18.15 Although aluminum has a negative standard potential, signifying that
it can be oxidized by hydrogen ions (as in the hydrochloric acid, left), nitric acid
(right) stops reacting with it as soon as an impenetrable layer of aluminum oxide
has formed on its surface. This resistance to further reaction is termed
passivation of the metal.
E 0 (Al3+ /Al)=-1.66 V
E (Al /Al)  E (H /H)  1.66 V
0
3+
0
+
Al(s)+H (aq)  Al (aq)+H 2 (g)
+
3+
There are no images in this section of the chapter.
Figure 18.16 The relation between the standard potential of a
reaction (reactants, purple; products, yellow) and the equilibrium
constant.
Investigating Matter 18.1 (a)
A glass electrode in a
protective plastic sleeve (left) is
used to measure pH. It is
sometimes used in conjunction
with a calomel electrode (right)
in pH meters such as this one.
Investigating Matter 18.1 (b) This durable portable pH meter can
be used for quick measurements of pH in the field. Its accuracy is
not as high as that of a laboratory pH meter.
Figure 18.17 A commercial dry cell
consists of a graphite cathode in a
zinc container; the latter acts as the
anode. The container is filled with a
moist paste of NH4Cl, MnO2, finely
divided carbon, and an inert filler
such as starch. In the cell reaction,
manganese(IV) is reduced to
manganese(III) and zinc metal is
oxidized to Zn2 ions.
Figure 18.18 One cell of a leadacid battery like those used in
automobiles. A lead-acid battery
is an example of a secondary cell.
It needs to be charged before it
can produce a current. The
electrolyte is dilute sulfuric acid.
Figure 18.19 A rechargeable
nickel-cadmium (nicad) cell. The
electrodes are assembled in a
jelly roll arrangement and
separated by a layer of paper
soaked in moist sodium or
potassium hydroxide.
Figure 18.20 The electric eel (Electrophorus electricus)
lives in the Amazon. The average potential difference it
produces along its length (1 m) is about 700 V.
Case Study 18 (a) One of the three alkali fuel cells used on the
space shuttle. Although only one cell is needed to provide lifesupport, electricity, and drinking water, shuttle flight rules require
that all three be functioning.
Case Study 18 (b) High-pressure hydrogen tanks run across the
top of this bus provided by Ballard Power Systems for testing
hydrogen-oxygen fuel cells in Chicago. The bus is pollution free
and can go 250 miles before needing to be refueled.
Figure 18.21 Iron nails stored
in oxygen-free water (left) do
not rust because the oxidizing
power of water itself is weak.
When oxygen is present (as a
result of air dissolving in the
water, right), oxidation is
thermodynamically
spontaneous and rust soon
forms.
Figure 18.22 The mechanism of rust formation. (a) Oxidation of
the iron occurs at a point out of contact with the oxygen of the air,
and the surface of the metal acts as an anode in a tiny galvanic
cell. (b) Further oxidation of Fe2 to Fe3 results in the deposition
of rust on the surface.
Figure 18.23 Metal girders are galvanized by immersion
in a bath of molten zinc.
Fig. 18.24 In the cathodic
protection of a buried
pipeline, or other large
metal construction, the
artifact is connected to a
number of buried blocks of
metal, such as magnesium
or zinc. The sacrificial
anodes (the magnesium
block in this illustration)
supply electrons to the
pipeline (the cathode of the
cell), thereby preserving it
from oxidation.
Figure 18.25 A schematic
picture of the electrolytic cell
used in the Dow process for
magnesium. The electrolyte is
molten magnesium chloride.
As the current generated by
the external source passes
through the cell, magnesium
metal is produced at the
cathode and chlorine gas is
produced at the anode.
Figure 18.26 In this schematic picture of an electrolysis experiment, the
electrons emerge from a galvanic cell at its anode () and enter the
electrolytic cell at its cathode (), where they bring about reduction.
Electrons are drawn out of the electrolytic cell through its anode () and
into the galvanic cell at its cathode (). If the cell reaction in the galvanic
cell is more strongly spontaneous than the reaction in the electrolytic
cell is nonspontaneous, then the overall process is spontaneous. This
experiment is an example of one reaction driving another to which it is
coupled.
Figure 18.27 Michael Faraday (1791–1867) giving a public
lecture on chemistry at the Royal Institution in London.
Figure 18.28 A schematic
picture showing the
electrolytic process for
refining copper. The anode
is impure copper. It
undergoes oxidation, and
the Cu2 ions so produced
migrate to the cathode,
where they are reduced to
pure copper metal. A similar
arrangement is used for
electroplating objects.
Figure 18.29 A schematic representation of the
stoichiometric relations that are used to calculate the
amount of product formed by electrolysis or the amount
of time current must flow to produce a given product.
Figure 18.30 In the Downs
process, molten sodium
chloride is electrolyzed with
a graphite anode (at which
the Cl ions are oxidized to
chlorine) and a steel
cathode (at which the Na
ions are reduced to
sodium). The sodium and
chlorine are kept apart by
the hoods surrounding the
electrodes. Calcium chloride
is present to lower the
melting point of sodium
chloride to a more
economical value.
Figure 18.31 Chromium plating lends decorative flair as well as
protection to the steel of this motorcycle. Large quantities of electricity
are needed to plate chromium because six electrons are required to
produce each atom of chromium.
Nobel Prize in Chemistry 2007
Assignment for Chapter 18
18.2; 18.6; 18.17; 18.27