Chapter 18 Electrochemistry
Principles of Chemistry
A Molecular Approach,
1 st Ed.
Nivaldo J. Tro
Dr. Azra Ghumman
Memorial University of Newfoundland
Electrochemistry
18.2
18.3
18.4
18.8
18.7
Balancing Oxidation-Reduction Reactions in Acidic and
Basic solution
Construction of Voltaic Cells, Notation for Voltaic Cells
Standard Cell emf’s and Standard Electrode potentials
Aqueous Electrolysis,Stoichiometry of Electrolysis
Batteries: Some Commercial Voltaic Cells
18.9
Corrosion: Undesirable Redox Reaction
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
2
18.2 Oxidation-Reduction Reactions
in Acidic and Basic Solutions
Oxidation-Reduction reactions always involve
transfer of electrons from one species to another.
– Oxidation is the loss of electrons (OIL)
– Reduction is the gain of electron (RIG)
• Recall that
– the species that loses electrons is oxidized,
– while the species that gains electrons is reduced.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
3
1
Reaction of Iron with Cu2+(aq)
• Fe(s) + Cu2+(aq)
Chem 1011
Fe2+(aq) + Cu(s)
Chapter 18 Electrochemistry
Dr. A. Ghumman
4
Oxidizing and Reducing agent
• An oxidizing agent is a species that oxidizes
another species; it is itself reduced by accepting e– A reactant in reduction half reaction
• A reducing agent is a species that reduces another
species; it is itself oxidized by losing e– A reactant in oxidation half reaction
Loss of 2 e- oxidation
• Reducing agent (Fe)
Fe( s ) Cu
oxidizing agent (Cu+2 )
Chem 1011
2
( aq )
Fe
Gain of 2
e-
2
( aq )
reduction
Cu( s )
Chapter 18 Electrochemistry
Dr. A. Ghumman
5
Oxidation-Reduction Reactions
• A cell is constructed to physically separate an
oxidation-reduction reaction into two halfreactions.
• Electrical current: Flow of electrons
– Redox reactions involve transfer of e- so they can
be used to generate electricity
• Voltage:The force with which electrons travel from
the oxidation half-reaction to the reduction halfreaction is measured as voltage.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
6
2
Electrochemistry
• Electrochemistry is the study of inter-conversion of
electrical and chemical energy and this interconversion takes place in an electrochemical cell.
• An electrochemical cell is a device in which a chemical
reaction(redox) either uses or generates an electric current.
– A voltaic, or galvanic cell (battery) is an
electrochemical cell in which a spontaneous reaction
generates an electric current.
– An electrolytic cell is an electrochemical cell in which
an electric current drives an otherwise nonspontaneous
reaction.
– we will discuss the basic principles behind these cells and
explore some of their commercial uses.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
7
18.2Construction of Voltaic Cells
•
•
•
A Voltaic cell consists of two half-cells that are electrically
connected.
Each half-cell is the portion of an electrochemical cell in which
a half reaction takes place;
– A zinc metal strip(Zn electrode) dipped into a solution of a
zinc salt is called a Zn-Zn2+ half cell,
– And a copper strip (Cu electrode) dipped into a solution of
copper salt is Cu2+-Cu half cell (Cu2+ solution is blue in
colour)
To produce electric current
1. Two half cells must be connected externally so that
electrons can flow from one metal electrode to another.
2. Two half cells must be connected internally by salt bridge to
allow the flow of ions between them.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
8
Voltaic Cell
• External circuit: In a voltaic cell, two half-cells are
connected externally in such a way that electrons
flow from one metal electrode to the other electrode.
• For example in Zn-Cu voltaic cell
– The electrons flow from Zn electrode through the
external circuit to the copper electrode where
copper ions gain the electrons to become copper
metal.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
9
3
A Zn-Cu Voltaic Cell
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
10
Voltaic Cell
• Salt Bridge (Internal circuit) is a inverted U shaped
tube that contains a strong electrolyte in a gel that
connects two half cells internally in a voltaic cell;
– allows the flow of ions but prevents the mixing of
the different solutions that would permit direct
reaction of the cell reactants.
• Why need salt Bridge?
– Without this internal connection, too much
positive charge builds up in the zinc half-cell (and
too much negative charge in the copper half-cell)
causing the reaction to stop.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
11
Voltaic Cell
• Anode:The electrode at which oxidation occurs is the anode
( -ve electrode)
– The electrons are lost, in the oxidation half-reaction.
Zn(s)
Zn2+( aq ) + 2e• Cathode: The electrode at which reduction occurs is called
cathode (+ve electrode)
– The electrons are gained, in the reduction half-reaction.
Cu(aq ) + 2eCu(s )
• Net cell reaction:The sum of two half reactions is the net
reaction that occur in the voltaic cell e.g.
Zn(s) + Cu2+(aq)
Chem 1011
Zn 2+(aq) + Cu(s)
Chapter 18 Electrochemistry
Dr. A. Ghumman
12
4
Summary of Voltaic Cell
Note that electrons are given up at the anode (Zn) and flow to
the cathode where they are used up in reduction reaction.
• Points to remember In a Voltaic cell;
• Electrons always flow from anode to cathode.
– Oxidation occurs at anode.
– Anode has a negative sign.
– Anions flow towards anode from salt bridge.
– Reduction occurs at cathode.
– Cathode has a positive sign.
– Cations flow towards cathode from salt bridge in a voltaic cell.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
13
Notation for Voltaic Cells
• Notation for Voltaic cell is a convenient way to designate a
particular voltaic cell, e.g. Zn-Cu cell is commonly abbreviated
as:
Zn (s ) | Zn 2 (aq ) || Cu 2 (aq ) | Cu (s )
anode
salt bridge
cathode
• The anode (oxidation half-cell) is always written on the left (in
the beginning)
• The cathode (reduction half-cell) is written on the right (at the
end)
• A single vertical line is a phase boundary between a solid and
an aqueous ion solution
• The double vertical line is used to indicate the salt bridge
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
14
Potential Difference
• Current is measured in amperes(A) 1A= 1C/s
• Electrical current is driven by a difference in potential
energy called Potential difference
• Electrode potential
• Potential difference is a measure of the difference
in potential energy(in joules) per unit
charge(coulomb)
– SI unit is Volt (V)
1 V = 1 J/C
• Cell potential (Ecell)or cell emf: The potential
difference between two electrodes in a voltaic cell
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
15
5
Notation for Voltaic Cells
•
•
Half- reaction involving a gas
– an inert material such as platinum serves as a terminal and an
electrode surface on which the reaction occurs.
– does not take part in the electrode reaction,
– only serves as a means of transferring e-s
Hydrogen electrode; Hydrogen bubbles over a platinum plate
immersed in an acidic solution.
– The cathode half-reaction is:
2H (aq ) 2e
•
H 2 (g )
Other cells may be written involving hydrogen where the half cell may
be either the anode or the cathode:
H2(g)
2H+ (aq) + 2e-
Chem 1011
oxidation (anode)
Chapter 18 Electrochemistry
Dr. A. Ghumman
16
Hydrogen Electrode
2H (aq ) 2e
H 2 (g )
•The notation for the hydrogen
electrode, written as a cathode, is:
H (aq ) | H 2 (g ) | Pt
•To write such an electrode as an anode,
you simply reverse the notation.
Pt | H 2 (g ) | H (aq )
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
17
Notation for Voltaic Cells
• To fully specify a voltaic cell, it is necessary to give
the concentration of solutions and the
pressures of gases.
• In the cell notation,these are written in parentheses
For example:
Zn(s) | Zn2 (1.0 M) || H (aq) | H2 (1.0 atm) | Pt
oxidation at anode
reduction at cathode
if multiple electrolytes are in same phase, a comma is used
rather than | (often use an inert electrode)
Fe(s) | Fe2+(aq) || MnO4 (aq), Mn2+(aq), H+ (aq) | Pt(s)
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
18
6
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
19
Writing the cell reaction from cell
notation
• Example:Give the overall cell reaction for the
voltaic cell
Cd ( s ) | Cd
2
( 1 . 0 M ) || H ( aq ) | H 2 ( 1 . 0 atm ) | Pt
Solution: Write the two half-cell reactions as:
Cd(s ) Cd 2 (aq ) 2e
2H (aq ) 2e
H 2 (g )
Add them together
Overall Rxn
Cd(s) 2H (aq)
Chem 1011
Cd2 (aq) H2 (g)
Chapter 18 Electrochemistry
Dr. A. Ghumman
20
Potential difference
• Potential difference: is the difference in electric
potential (electrical pressure) between two electrodes.
• Electrode potential:
• SI unit of Potential: The volt, (V) is the SI unit of
potential difference equivalent to 1 joule of energy
per coulomb of charge.
1 volt
1J
C
• Two other terms to describe potential difference are, cell
potential Ecell and Electromotive force emf of a cell.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
21
7
Electromotive Force
• The movement of electrons is analogous to the pumping of
water from one point to another.
• An electric charge moves from a point of high electrical
potential (high PE) to one of lower electrical potential (lower
PE).
– e.g. An e- in Zn has higher potential energy than an e- in Cu
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
22
Standard Electrode Potentials
• A cell emf is a measure of the driving force of the
cell reaction.
• The reaction at the anode has a definite oxidation
potential;
Reduced species
Oxidized species +n e - ( oxi)
while the reaction at the cathode has a definite
reduction potential;
Oxidized species + n e-
Reduced species (red)
• Thus, the overall Ecell is the difference between these
two potentials.
Ecell = Ecathode - E anode
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
23
Measuring ECell
Ecell:The maximum potential difference between
the electrodes of a voltaic cell
• It can be measured by an electronic digital
voltmeter which draws negligible current
• When anode is connected to –ve terminal
and cathode is connected to + ve terminal
of the voltmeter, you will get a +ve voltage.
• The Ecell is an intensive property,
– independent of the number of e- passing through the cell
– Dependent on the nature of the redox reaction and the
concentrations of the species involved
Ecell = Ecathode - E anode
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
24
8
Reduction Potential
• A reduction potential is a measure of the tendency
to gain electrons in the reduction half-reaction
– You can look at the oxidation half-reaction as the
reverse of a corresponding reduction reaction.
• The oxidation potential is negative of the
reduction potential for the reverse reaction.
Eoxi = - Ered
• By convention we tabulate reduction potential and call
them electrode potentials (E).
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
25
• Standard potentials for oxidation half reactions are obtained by
reversing the sign of the standard reduction potentials
• For a general reaction
X
X+ + eEoxi = - Ered
• Listed reduction potential for the following metals and
their corresponding oxidation potentials
Fe2+ (aq) +2eFe(s) Ered = -409 V
Fe(s)
Fe2+ (aq) +2e- Eoxi = + 409 V
2+
Cu (aq) +2e
Cu(s) Ered = + 0.339 V
Cu(s)
Cu2+ (aq) +2e- Eoxi = - 0.339V
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
26
Standard Cell potential
•.E°cell is the standard potential of a cell operating under standard
conditions.
• standard conditions: solute concentration of 1 M, and gas
pressures of 1 atm at 25°C.
– It is not possible to measure the single electrode potentials.
Only E° cell can be measured, so we choose a reference
electrode
• Standard H2 Electrode(SHE): By convention the reference
electrode is SHE
– Electrode potential of the SHE is exactly zero
– Standard electrode potentials are measured relative to this
hydrogen reference electrode.
2H+ (aq 1M) + 2eH2(g 1 atm) E0red = 0.00V
H2(g 1 bar)
2H+ (aq 1M) + 2e- E0oxi = 0.00V
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
27
9
Potential Energy and the Voltage
The more –ve the electrode potential, the greater the
PE of an electron at that electrode because –ve charge
repels electrons
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
28
Standard Cell Potential
• For example, the emf of a cell composed of a zinc
electrode connected to a hydrogen electrode,
Red. half 2H+ (aq) + 2eH2(g)
Oxi. half Zn(s)
Zn2+(aq) + 2eH2(g) + Zn2+(aq)
E° = 0.00V
E° = –0.76 V
2H+ (aq)+ Zn(s)
E°cell= 0.76 V
Ecell = E H2 + (-E Zn)
= 0.76 V
– Since zinc acts as the anode in this cell, its
reduction potential is listed as –0.76 V.
• Listed potentials in Table 18.1 are standard electrode
potentials for the reduction half reactions
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
29
Standard Electrode Potentials(Table
18.1)
Strongest
oxidizing agent
at the top
(largest E° + ve
value)
Strongest
reducing agent
at the bottom
(smallest E°
-ve value)
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
30
10
Table 18.1 Cont’d
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
31
Strength of Oxidizing
• Electrode potentials are useful in determining the strengths of
oxidizing and reducing agents under standard-state
conditions
• Oxidizing agent is a species that can gain electron easily.
– is a reactant in reduction half reaction
– has a strong attraction for the e- and hence, can easily
oxidize other species e.g. F2(g)
– Largest the value of E red ( most positive), the strongest
the oxidizing agent is (at the top of the table).
F2 (g) + 2e2F-(aq)
E° = 2.87 V
• H+ ion is fairly strong oxidizing agent and oxidize many metals
including Zn and Mg.
• Common strong oxidizing agent used in the lab Cr2O7 2- MnO4 2and Cl2
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
32
Strength of Reducing agent
• Reducing agent is a species that can lose electron
easily (bottom of the table).
– The strongest reducing agents in a table of
standard electrode potentials are the reduced
species corresponding to the half-reactions with
the smallest ( most negative) Eored values e.g.
Li(s), Ca(s) bottom of the table)
– Strength of reducing agent increase from top to
the bottom of the table (strongest at the bottom).
Li+ (aq) + eChem 1011
Li(s)
E° = -3.04 V
Chapter 18 Electrochemistry
Dr. A. Ghumman
33
11
Determining the Relative strength of
Oxidizing and Reducing agents
Example: a. Arrange the following oxidizing agents in order of
increasing strength under standard conditions.
Br2(l), Fe3+(aq) and Cr2O72-.
b. Arrange the following reducing agents in order of increasing
strength under standard conditions.
Al(s), Na(s), and Zn(s)
Points to remember
Strongest oxidizing agents, largest value of Eored (top of the
table)
Strongest reducing agents, smallest value of Eored(bottom
of the tabl)
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
34
Calculating Cell emf’s from Standard
Potentials
• Standard voltage for an oxidation-reduction reaction is the sum
of the standard voltages for two half reactions; the oxidation
half and reduction half.
• It can be measured experimentally and
• It can be calculated from tabulated standard reduction potential
values.
1. Combining half reactions and their standard potential
values (reduction and oxidation).
2. Using an equation to calculate the difference between
standard reduction potentials i.e.
• E cell =E red - E oxid (when both are written as reductions)
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
35
Calculating Cell’s emf from Standard
Potentials by combining half
reaction
• Points to remember about cell voltage
– The calculated E cell is always a positive
quantity for a reaction taking place in a voltaic
cell.
• A negative cell emf indicates the cell reaction is
non-spontaneous.
– The quantities E cell, E oxi, and E red are
independent of how the equation for the cell
reaction is written
When you reverse the equation, change the sign for E°.
You never multiply the voltage values by the
coefficients of the balanced equation.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
36
12
Calculating Cell emf’s from Standard
Potentials by combining half reaction
• Calculate E cell for the following reaction at 25 C:
•
IO3–(aq) + 6 H+ (aq) + 5 I- (aq)
3 I2(s) + 3 H2O(l)
Separate the reaction into the oxidation and reduction halfreactions.
ox (anode):
Red (cathode):
2 I- (s)
I 2(aq) + 2 eIO3- (aq) + 6 H+ (aq) + 5 e-
½ I 2(s) + 3 H2O(l)
Find the E for each half-reaction and sum to get E cell.
E
E cell
- = -E
red of I2 = - 0.54 V
E red of IO3- = +1.20 V
= (- 0.54 V) + (+1.20 V) = +0.66 V
ox of I
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
37
Spontaneity of Redox Reactions
Spontaneity of Redox Reactions
• – Redox reactions are spontaneous if the cell
potential is a positive quantity.
• – The cell reaction is non spontaneous if the cell
potential is a negative value.
• In general any reduction half-reaction will be
spontaneous when paired with the reverse of a half –
reaction(oxidation) that appears below it in the table.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
38
Spontaneity of Redox Reactions
• Example: Predict whether the following reaction is
spontaneous under standard conditions:
Fe(s) + Mg2+(aq)
Fe2+(aq) + Mg(s).
• Solution: Separate into two half-reactions
Fe2+(aq) + 2 e-
oxidation:
Fe(s)
reduction:
Mg2+(aq) + 2 e-
Mg(s)
• Reverse the first half and look up the relative position
of reduction half position in the table.
Since Mg2+ reduction is below Fe2+ reduction, the
reaction is NOT spontaneous as written.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
39
13
Predicting whether a metal will
dissolve in Acid
• In general metals whose reduction half reactions lie
below the reduction of H+ to H2 in Table 18.1 will
dissolve in acids while metals above it will not.
• Almost all metals will dissolve in HNO3.
– having N reduced rather than H
– Au and Pt dissolve in HNO3 + HCl.
NO3- (aq) + 4 H+(aq) + 3 eH2O(l)
Chem 1011
NO(g) + 2
Chapter 18 Electrochemistry
Dr. A. Ghumman
40
Electrolysis
• An electrolytic cell is an electrochemical cell in which an
electric current drives an otherwise non-spontaneous reaction.
– A non spontaneous cell reaction can be forced to
proceed by pumping e- into the system
• The following can be forced to proceed in the direction indicated
by applying an external voltage larger than +0.748V.)
Cu(s) + Fe2+( aq)
Cu2+( aq) + Fe(s)
Ecell = -0.748V
• If the applied voltage is equal equal to +0.748V the system will
be at equilibrium
The reaction Overvoltages: In real life, a higher applied voltage is
required to overcome cell resistance, and any change due to
non standard conditions.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
41
Comparison of Voltaic and
Electrolytic cell
– In all electrochemical cells oxidation occurs at anode and
reduction occurs a cathode.
• In Voltaic Cells:
– The anode is the source of electrons and has a negative
charge( anode -ve)
– The cathode draws electrons and has a positive
charge(cathode +ve)
• In Electrolytic cells:
– anode is +ve and cathode is –ve,
– e- are pulled from anode by the +ve terminal of the battery
– and e- are pumped into cathode from the negative terminal
of battery
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
42
14
Comparison of Voltaic and
Electrolytic cell
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
43
Applications of Electrolysis
• Electrolysis: The process of forcing a chemical reaction by
applying current in an electrolytic cell.
– Many important substances,like, Na, Al metal and Cl2(g) are
produced by electrolysis of their molten salts. Li, Mg, and
Ca, metals are all obtained by the electrolysis of their
chlorides.
• A Downs cell is a commercial electrochemical cell used to
obtain sodium metal by electrolysis of molten NaCl.
• Na forms at cathode, from the reduction of Na+ ions.
• Cl2(g) forms at the anode from the oxidation of Cl- ions.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
44
Electrolysis
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
45
15
Electroplating
• In electroplating, the object to be
plated is the cathode.
• Cations are reduced at the
cathode and plate to the surface of
the metal object.
• The anode is made of the plate
metal. The anode oxidizes and
replaces the metal cations in the
solution.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
46
Stoichiometry of Electrolysis
Involve the measurement of numbers of electrons.
– You do not weigh electrons.
– Rather, you measure the quantity of electric
charge that has passed through a circuit.
– To determine this we must know the current
and the length of time it has been flowing for.
• Faraday Law: The quantity of a substance
deposited at an electrode is directly proportional to
the quantity of electric charge that had passed in the
circuit.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
47
Stoichiometry of Electrolysis
• Electric current is measured in amperes.
• An ampere (A) is the base SI unit of current
equivalent to 1 coulomb/second (C/s).
– ‘The quantity of electric charge passing through a
circuit in a given amount of time is given by:
Electric charge (C) = electric current (A) time lapse (s)
Coulomb C = Ampere (C/s) x t(s)
• The Faraday constant, (F ) is the magnitude of
charge on one mole of electrons; it equals 96,500
coulombs (9.65 x 104 C).
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
48
16
Stoichiometry of Electrolysis
•
To determine the # of mole of e- involved in
electrolysis reaction we can write,
# mol of e- = current (C/s) x time(s) x 1mol e96500C
• To determine the mass of a product in an electrolysis
reaction,we can follow the following pathway:
Current and time
Coulombs
mol emol of
product
grams of product.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
49
Stoichiometry of Electrolysis
• Example: When an aqueous solution of potassium
iodide is electrolyzed using platinum electrodes, the
half-reactions are:
2I (aq )
2H 2O(l ) 2e
I 2 (aq ) 2e
H 2 (g ) 2OH (aq )
– How many grams of iodine are produced when a
current of 8.52 mA flows through the cell for 10.0
min?
– When the current flows for 6.00 x 102 s (10.0
min), the amount of charge is:
(8.52 10 3 A ) (6.00 102 s ) 5.11 C
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
50
Cont’d
– Note that two moles of electrons are equivalent to one
mole of I2. Hence,
5 .11 C
1 mol e
9 .65 10 4 C
1 mol I 2
2 mol e
254 g I 2
1 mol I 2
6 .73 10
3
g I2
• Example:In commercial preparation of aluminum, Al2O3
is electrolyzed at 1000°C. Assume the cathode reaction is
Al3+(aq) + 3eAl(s)
• How many coulombs of electricity are required to give
2.78 kg of Al?
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
51
17
Some Commercial Voltaic Cells
• The Leclanché dry cell, or zinc-carbon dry cell, is a
voltaic cell with a zinc can as the anode and a
graphite rod in the center surrounded by a paste of
manganese dioxide, ammonium and zinc chlorides,
and carbon black as the cathode.
– The electrode reactions are:
Zn(s )
Zn 2 (aq ) 2e
2NH 4 (aq) 2MnO 2 (s ) 2e
anode
Mn 2O 3 (s ) H 2O( l ) 2NH 3 (aq )
cathode
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
52
Leclanché Dry Cell
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
53
• The alkaline dry cell is similar to the Leclanché cell,
but it has potassium hydroxide in place of ammonium
chloride.
– The electrode reactions are:
Zn(s )
Zn 2 (aq ) 2e
MnO 2 (s ) H 2O( l ) 2e
Chem 1011
Mn 2O 3 (s ) 2OH (aq )
Chapter 18 Electrochemistry
Dr. A. Ghumman
anode
cathode
54
18
A small alkaline dry cell
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
55
A lead storage battery
• The lead storage cell (a rechargeable cell) consists of
electrodes of lead alloy grids; one electrode is packed with a
spongy lead to form the anode, and the other electrode is
packed with lead dioxide to form the cathode.
• six cells in series
• electrolyte = 30% H2SO4
The electrode reactions are
Anode: Pb
Pb(s) + SO42- (aq)
PbSO4(s) + 2 e-
Cathode: Pb coated with PbO2 is cathode and PbO2 is reduce
PbO2(s) + 4 H+ (aq) + SO42- (aq) + 2 ePbSO4(s) + 2 H2O(l)
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
56
A lead storage battery
• six cells in series
• electrolyte = 30% H2SO4
•Anode:
Pb(s) + SO42- (aq)
PbSO4(s) + 2 e-
•Cathode: Pb coated with PbO2
•PbO2 is reduced
cell voltage = 2.09 V
•rechargeable, heavy
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
57
19
NiCad storage batteries
• The nickel-cadmium cell (nicad cell) consists of an
anode of cadmium and a cathode of hydrated nickel
oxide on nickel; the electrolyte is potassium
hydroxide.
– The electrode reactions are:
Cd(s ) 2OH (aq )
NiOOH(s ) H 2O(l ) e
Chem 1011
Cd(OH )2 (s ) 2e
anode
Ni(OH )2 (s ) OH (aq ) cathode
Chapter 18 Electrochemistry
Dr. A. Ghumman
58
Nicad storage batteries
Photo courtesy of American Color.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
59
Fuel Cell
• A fuel cell is essentially a battery, but differs by
operating with a continuous supply of energetic
reactants, or fuel.
• The first fuel-cell was invented in 1839 by William
Groves
– For a hydrogen-oxygen fuel cell, the electrode
reactions are:
2 H 2 ( g ) 4OH (aq )
O 2 ( g ) 2 H 2 O ( l ) 4e
H 2O( l ) 4e
4OH (aq )
anode
cathode
• MFC, DEFC and Formic acid Fuel Cells
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
60
20
Fuel-Cell
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
61
Electrochemical Process involved in
Rusting
• Rust is hydrated iron(III) oxide.
• Moisture must be present.
Water is a reactant.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
62
Corrosion and its Prevention
• Corrosion is the spontaneous oxidation of a metal
by chemicals in the environment.
• Since many materials we use are active metals,
corrosion can be a very big problem.
• Rusting can be prevented by placing a sacrificial
electrode in electrical contact with iron.
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
63
21
Operational Skills
• Balancing equations in acidic and basic solutions by
the half-reaction method
• Writing the cell reaction from the cell notation
• Determining the relative strengths of oxidizing and
reducing agents
• Determining the direction of spontaneity from
electrode potentials
• Calculating the emf from standard potentials
• Relating the amounts of product and charge in an
electrolysis
Chem 1011
Chapter 18 Electrochemistry
Dr. A. Ghumman
64
22
This document was created with Win2PDF available at http://www.win2pdf.com.
The unregistered version of Win2PDF is for evaluation or non-commercial use only.
This page will not be added after purchasing Win2PDF.