C h a pt e r
18
Electrochemistry
Chemistry 4th Edition
McMurry/Fay
Dr. Paul Charlesworth
Michigan Technological University
Redox Reactions
Prentice Hall ©2004
Chapter18
Redox Reactions
01
Slide 2
01
Redox reaction are those involving the oxidation
and reduction of species.
OIL – Oxidation Is Loss of electrons.
RIG – Reduction Is Gain of electrons.
Oxidation and reduction must occur together.
They cannot exist alone.
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Chapter18
Slide 3
1
Redox Reactions
Prentice Hall ©2004
Chapter18
Redox Reactions
02
Slide 4
02
Oxidation Half-Reaction: Zn(s) → Zn2+(aq) + 2 e–.
• The Zn loses two electrons to form Zn2+.
•
Prentice Hall ©2004
Chapter18
Redox Reactions
•
•
Slide 5
03
Reduction Half-Reaction: Cu2+(aq) + 2 e– → Cu(s)
The Cu 2+ gains two electrons to form copper.
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Chapter18
Slide 6
2
Redox Reactions
•
04
Overall: Zn(s) + Cu 2+(aq) → Zn2+(aq) + Cu(s)
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Chapter18
Electrochemical Cells
Slide 7
01
•
Electrodes: are usually metal strips/wires
connected by an electrically conducting wire.
•
Salt Bridge: is a U-shaped tube that contains a gel
permeated with a solution of an inert electrolyte.
•
Anode: is the electrode where oxidation takes
place.
•
Cathode: is the electrode where reduction takes
place.
Prentice Hall ©2004
Chapter18
Electrochemical Cells
•
Slide 8
02
Convention for expressing the cell:
Anode Half-Cell || Cathode Half-Cell
Electrode | Anode Soln || Cathode Soln | Electrode
Zn(s) | Zn 2+ (1 M) || Cu 2+ (1 M) | Cu(s)
Pt(s) | H2 (1 atm) | H+ (1 M) || Fe 3+(aq), Fe 2+(aq) | Pt(s)
•
Electrons flow from anode to cathode. Anode is
placed on left by convention.
Prentice Hall ©2004
Chapter18
Slide 9
3
Electrochemical Cells
•
03
The standard potential of any galvanic cell is the
sum of the standard half-cell potentials for the
oxidation and reduction half-cells.
E°cell = E°oxidation + E°reduction
•
Standard half-cell potentials are always quoted as
a reduction process. The sign must be changed for
the oxidation process.
Prentice Hall ©2004
Chapter18
Electrochemical Cells
Slide 10
04
•
The standard half-cell potentials are determined
from the difference between two electrodes.
•
The reference point is called the standard hydrogen
electrode (S.H.E.) and consists of a platinum
electrode in contact with H2 gas (1 atm) and
aqueous H+ ions (1 M).
•
The standard hydrogen electrode is assigned an
arbitrary value of exactly 0.00 V.
Prentice Hall ©2004
Chapter18
Electrochemical Cells
Prentice Hall ©2004
Chapter18
Slide 11
05
Slide 12
4
Electrochemical Cells
Prentice Hall ©2004
Chapter18
Electrochemical Cells
•
When selecting two half-cell reactions the more
negative value will form the oxidation half-cell.
•
Consider the reaction between zinc and silver:
•
Ag+ (aq) + e– → Ag(s)
E° = 0.80 V
Zn2+ (aq) + 2 e– → Zn(s)
E° = – 0.76 V
06
Slide 13
07
Therefore, zinc forms the oxidation half-cell:
Zn(s) → Zn2+ (aq) + 2 e–
Prentice Hall ©2004
E° = – (–0.76 V)
Chapter18
Electrochemical Cells
Slide 14
08
•
What is the standard emf of an electrochemical cell
made of a Cd electrode in a 1.0 M Cd(NO3) 2
solution and a Cr electrode in a 1.0 M Cr(NO3) 2
solution?
•
What is the standard emf of an electrochemical cell
made of a Mg electrode in a 1.0 M Mg(NO3) 2
solution and a Ag electrode in a 1.0 M AgNO3
solution?
Prentice Hall ©2004
Chapter18
Slide 15
5
Spontaneity of a Reaction
01
•
The value of E°cell is related to the thermodynamic
quantities of ?G° and K.
•
The value of E°cell is related to ?G° by:
?G° = –nFE°cell
•
The value of K is related to ?G° by:
?G° = –RT ln K
Prentice Hall ©2004
Chapter18
Spontaneity of a Reaction
•
Slide 16
02
These are combined:
–nFE°cell = –RT ln K
•
and rearranged at 25°C to give:
or
•
E°cell = (0.0257/n) ln K
E°cell = (0.0592/n) log K
(Units = volts, n = moles, F = 96,500 coulombs)
Prentice Hall ©2004
Chapter18
Spontaneity of a Reaction
•
Slide 17
03
Calculate the standard free energy change (?G°)
and the equilibrium constant (K) for the following
reactions at 25°C:
1. Sn(s) + 2 Cu 2+(aq) æ Sn2+(aq) + 2 Cu +(aq)
2. Fe2+(aq) + 2 Ag(s) æ Fe(s) + 2 Ag+(aq)
3. 4 Fe 2+( aq) + O2(g) + 4 H+(a q) → 4 Fe 3+( aq) + 2 H2 O(l)
Prentice Hall ©2004
Chapter18
Slide 18
6
The Nernst Equation
•
01
Cell potentials can be modified by temperature and
composition changes according to the equation:
?G = ?G° + RT lnQ
∴ –nFE = –nFE° + RT lnQ
•
Giving:
Prentice Hall ©2004
E = Eo −
RT
ln Q
nF
Chapter18
The Nernst Equation
Slide 19
03
•
Consider the reaction of metallic zinc with
hydrochloric acid. Calculate the cell potential at
25°C when [H+] = 1.0 M, [Zn2 +] = 0.0010 M, and
PH2 = 0.10 atm.
•
Consider the reaction of metallic copper with
iron(III) to give copper(II) and iron(II). What is the
potential of a cell when [Fe 3+] = 0.0001 M, [Cu 2 +] =
0.25 M, and [Fe 2+] = 0.20 M?
Prentice Hall ©2004
Chapter18
The Nernst Equation
•
Slide 20
04
A particularly important use of the Nernst equation
is in the electrochemical determination of pH.
Pt | H2 (1 atm) | H+ (? M) || Reference Cathode
•
Ecell = EH2 → H+ + Eref
The Nernst equation can be applied to the
half-reaction:
H2(g) → 2 H+(aq) + 2 e–
Prentice Hall ©2004
Chapter18
Slide 21
7
The Nernst Equation
•
For the half-reaction: H2(g) → 2 H+(aq) + 2 e–
E H 2 → H + = E oH2
•
05
→ H+
+
V log H
– 0.0592
n
PH
2
2
E° = 0 V for this reaction (standard hydrogen
electrode). PH2 is 1 atm.
V log H+ 2
E H 2 → H+ = 0 V – 0.0592
n
Prentice Hall ©2004
Chapter18
Slide 22
The Nernst Equation
•
07
The overall potential is given by:
Ecell = (0.0592 V )(pH) + Eref
•
Which rearranges to give an equation for the
determination of pH:
Ecell − Eref
= pH
0.0592 V
Prentice Hall ©2004
Chapter18
The Nernst Equation
Slide 23
08
•
A higher cell potential indicates a higher pH,
therefore we can measure pH by measuring Ecell .
•
A glass electrode (Ag/AgCl wire in dilute HCl) with
a calomel reference is the most common
arrangement.
Glass: Ag(s) + Cl–(aq) → AgCl(s) + e–
E° = –0.22 V
Calomel: Hg2Cl2(s) + 2 e– → 2 Hg(l) + 2 Cl–(aq)
E° = 0.28 V
Prentice Hall ©2004
Chapter18
Slide 24
8
The Nernst Equation
09
• The glass pH probe is constructed as follows:
Ag(s) | AgCl(s) | HCl(aq) | glass | H+(aq) || reference
•The difference in [H+] from one side of the glass
membrane to the other causes a potential to develop,
which adds to the measured Ecell .
Ecell − Eref
= pH
0.0592 V
Prentice Hall ©2004
Chapter18
The Nernst Equation
•
Slide 25
10
The following cell has a potential of 0.55 V at 25°C:
Pt(s) | H 2 (1 atm) | H+ (? M) || Cl– (1 M) | Hg2Cl2(s) | Hg(l)
•
What is the pH of the solution at the anode?
•
The following cell has a potential of 0.28 V at 25°C:
Pt(s) | H 2 (1 atm) | H+ (? M) || Pb 2+ (1 M) | Pb(s)
•
What is the pH of the solution at the anode?
Prentice Hall ©2004
Chapter18
Batteries
Slide 26
01
•
Batteries are the most important practical
application of galvanic cells.
•
Single-cell batteries consist of one galvanic cell.
•
Multicell batteries consist of several galvanic cells
linked in series to obtain the desired voltage.
Prentice Hall ©2004
Chapter18
Slide 27
9
Batteries
02
Lead Storage Battery: A typical 12 volt battery
consists of six individual cells connected in series.
•
Anode: Lead grid packed with spongy lead.
•
Pb(s) + HSO4–(aq) → PbSO4(s) + H+ (aq) + 2 e–
Cathode: Lead grid packed with lead oxide.
PbO2(s) + 3 H + (aq) + HSO 4–(aq) + 2 e– → PbSO 4(s) + 2 H2O(l)
Electrolyte: 38% by mass sulfuric acid.
Cell Potential: 1.924 V
Prentice Hall ©2004
Chapter18
Batteries
•
Slide 28
03
Zinc Dry-Cell: Also called a Leclanché cell, uses a
viscous paste rather than a liquid solution.
Anode: Zinc metal can on outside of cell.
Zn(s) → Zn2+ (aq) + 2 e–
•
Cathode: MnO 2 and carbon black paste on graphite.
2 MnO2(s) + 2 NH 4+ (aq) + 2 e– → Mn2O3(s) + 2 NH 3(aq) +
2 H 2O(l)
•
•
Electrolyte: NH4Cl and ZnCl2 paste.
•
Cell Potential: 1.5 V but deteriorates to 0.8 V with use.
Prentice Hall ©2004
Chapter18
Batteries
•
Slide 29
04
Alkaline Dry-Cell: Modified Leclanché cell which
replaces NH4 Cl with NaOH or KOH.
Anode: Zinc metal can on outside of cell.
Zn(s) + 2 OH –(aq) → ZnO(s) + H2O(l) + 2 e–
Cathode: MnO 2 and carbon black paste on graphite.
2 MnO2 (s) + H2O(l) + 2 e– → Mn 2O3(s) + 2 OH –(aq )
Electrolyte: NaOH or KOH, and Zn(OH)2 paste.
Cell Potential: 1.5 V but longer lasting, higher power,
and more stable current and voltage.
Prentice Hall ©2004
Chapter18
Slide 30
10
Batteries
•
05
Mercury Dry-Cell: Modified Leclanché cell which
replaces MnO2 with HgO and uses a steel cathode.
Anode: Zinc metal can on outside of cell.
Zn(s) + 2 OH –(aq) → ZnO(s) + H2O(l) + 2 e–
Cathode: HgO in contact with steel.
2 HgO (s) + H2O(l) + 2 e– → Hg(l) + 2 OH –(aq)
Electrolyte: KOH, and Zn(OH)2 paste.
Cell Potential: 1.3 V with small size, longer lasting,
and stable current and voltage.
Prentice Hall ©2004
Chapter18
Batteries
•
Slide 31
06
Nickel–Cadmium Battery: Modified Leclanché cell
which is rechargeable.
Anode: Cadmium metal.
Cd(s) + 2 OH –(aq) → Cd(OH)2(s) + 2 e–
Cathode: Nickel(III) compound on nickel metal.
NiO(OH) (s) + H2O(l) + e– → Ni(OH)2(s) + OH –(aq)
Electrolyte: Nickel oxyhydroxide, NiO(OH).
Cell Potential: 1.30 V
Prentice Hall ©2004
Chapter18
Batteries
•
•
Slide 32
07
Nickel –Metal–Hydride
(NiMH ):
Replaces toxic Cd anode with
a hydrogen atom impregnated
ZrNi2 metal alloy.
•
During oxidation at the anode,
hydrogen atoms are released
as H 2O.
•
Recharging reverses this
reaction.
Prentice Hall ©2004
Chapter18
Slide 33
11
Batteries
•
08
Lithium Ion (Li–ion): The newest rechargeable
battery is based on the migration of Li+ ions.
•Anode:
Li metal, or Li atom impregnated graphite.
Li(s) → Li+ + e–
Cathode: Metal oxide or sulfide that can accept Li +.
MnO 2(s) + Li+ (aq) + e– → LiMnO2(s)
Electrolyte: Lithium-containing salt such as LiClO4, in
organic solvent. Solid state polymers can also be used.
Cell Potential: 3.0 V
Prentice Hall ©2004
Chapter18
Batteries
•
Slide 34
09
Fuel Cell: Uses externally fed CH4 or H2, which
react to form water. Most common is H2.
•
Anode: Porous carbon containing metallic catalysts.
2 H2(s) + 4 OH –(aq) → 4 H2O(l) + 4 e–
Cathode: Porous carbon containing metallic catalysts.
O2(s) + 2 H2O(l) + 4 e– → 4 OH –(aq)
Electrolyte: Hot aqueous KOH solution.
Cell Potential: 1.23 V, but is only 40% of cell capacity.
Prentice Hall ©2004
Chapter18
Batteries
10
•
Fuel cells are not batteries
because they are not self contained.
•
Fuel cells typically have
about 40% conversion to
electricity; the remainder is
lost as heat.
•
Excess heat can be used to
drive turbine generators.
Prentice Hall ©2004
Slide 35
Chapter18
Slide 36
12
Corrosion
01
•
Corrosion is the oxidative deterioration of metal.
•
25% of steel produced in USA goes to replace steel
structures and products destroyed by corrosion.
•
Rusting of iron requires the presence of BOTH
oxygen and water.
•
Rusting results from tiny galvanic cells formed by
water droplets.
Prentice Hall ©2004
Chapter18
Corrosion
Slide 37
02
Oxidation: Fe(s) → Fe2+ (aq) + 2 e–
Reduction: O2(g) + 4 H+ (aq) + 4 e– → 2 H2O(l)
Overall:
2 Fe(s) + O2(g) + 4 H+ (aq) → 2 Fe2+ (aq) + 2 H2O(l)
Prentice Hall ©2004
Chapter18
Corrosion
Slide 38
03
•
Galvanizing: is the coating of iron with zinc. Zinc is
more easily oxidized than iron, which protects and
reverses oxidation of the iron.
•
Cathodic Protection: is the protection of a metal
from corrosion by connecting it to a metal (a
sacrificial anode) that is more easily oxidized.
•
All that is required is an electrical connection to the
sacrificial anode (usually magnesium or zinc).
Prentice Hall ©2004
Chapter18
Slide 39
13
Corrosion
•
04
Cathodic Protection with Zinc Layer
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Chapter18
Corrosion
•
Slide 40
05
Cathodic Protection with Magnesium Anode
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Chapter18
Electrolysis
Prentice Hall ©2004
Slide 41
01
Chapter18
Slide 42
14
Electrolysis
01
•
Electrolysis: is the process in which electrical
energy is used to drive a nonspontaneous chemical
reaction.
•
An electrolytic cell is an apparatus for carrying out
electrolysis.
•
Processes in an electrolytic cell are the reverse of
those in a galvanic cell.
Prentice Hall ©2004
Chapter18
Electrolysis
Slide 43
05
•
Electrolysis of Water: Requires an electrolyte
species, that is less easily oxidized and reduced
than water, to carry the current.
•
Anode: Water is oxidized to oxygen gas.
2 H2O(l) → O2(g) + 4 H+(aq) + 4 e–
•
Cathode: Water is reduced to hydrogen gas.
4 H2O(l) + 4 e– → 2 H2(g) + 4 OH–(aq)
Prentice Hall ©2004
Chapter18
Electrolysis
Slide 44
02
•Electrolysis
ElectrolysisofofAqueous
Molten Sodium
SodiumChloride:
Chloride: uses
different processes to molten sodium chloride.
Based on cell potentials, water (–0.83 V) would be
preferentially reduced over sodium ions (–2.71 V).
Based on cell potentials, water (+1.23 V) should be
preferentially oxidized over chloride ions (+1.36 V).
The observed product at the anode is Cl2, not O2,
because of a phenomenon called overvoltage.
Prentice Hall ©2004
Chapter18
Slide 45
15
Electrolysis
04
•
Overvoltage : Additional voltage needed to
maintain the rate of electron transfer at the
electrode–solution interface.
•
Overvoltage is required when a half-reaction has a
significant activation energy, and so a slow rate.
•
Overvoltage for formation of O2 or H2 is much
greater than for formation of Cl2.
Prentice Hall ©2004
Chapter18
Electrolysis
Slide 46
07
•
Quantitative Electrolysis:The amount of
substance produced at an electrode by
electrolysis depends on the quantity of charge
passed through the cell.
•
Reduction of 1 mol of sodium ions requires 1 mol
of electrons to pass through the system.
•
The charge on 1 mol of electrons is 96,500
coulombs.
Prentice Hall ©2004
Chapter18
Electrolysis
Slide 47
08
•
To determine the moles of electrons passed, we
measure the current and time that the current flows:
•
Because the charge on 1 mol of e– is 96,500 C, the
number of moles of e– passed through the cell is:
Charge (C) = Current (A) x Time (s)
1 mole e –
Moles of e – = Charge (C) ×
96,500 C
Prentice Hall ©2004
Chapter18
Slide 48
16
Electrolysis
09
•
A constant current of 30.0 A is passed through an
aqueous solution of NaCl for a time of 1.00 hour.
How many grams of Na and how many liters of Cl2
gas at STP are produced?
•
How many amperes must be passed through a
Downs cell to produce sodium metal at a rate of
30.0 kg/hour?
Prentice Hall ©2004
Chapter18
Electrolysis Applications
•
01
Manufacture of Sodium (Downs Cell):
Prentice Hall ©2004
Chapter18
Electrolysis Applications
•
Slide 49
Slide 50
02
Manufacture of Cl2 and NaOH (Chlor–Alkali):
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Chapter18
Slide 51
17
Electrolysis Applications
•
Manufacture of Aluminum (Hall–Heroult):
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Chapter18
Electrolysis Applications
•
03
Slide 52
04
Electrorefining and Electroplating:
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Chapter18
Slide 53
18