Lecture Presentation
CHEM 111
Electrochemistry
Electrochemistry
• Redox reactions
• Electrochemical cells
• Electrode processes
• Construction
• Notation
• Cell potential and Go
• Standard reduction potentials (Eo)
• Non-equilibrium conditions (Q)
• Batteries
• corrosion
Electric
automobile
Electrochemistry
Electrochemical processes: Oxidation-Reduction
reactions in which:
• The chemical energy released by a spontaneous
reaction is converted to electrical energy or
• Electrical energy is used to cause a nonspontaneous
reaction to occur (or to trigger a chemical
change/reaction).
0
0
2+ 2-
2Mg (s) + O2 (g)
2Mg
O2 + 4e-
2MgO (s)
2Mg2+ + 4e- Oxidation half-reaction (lose e-)
2O2-
Reduction half-reaction (gain e-)
3
Electrochemistry
 Electrochemical cell: A system consisting of
electrodes and electrolytes where a chemical
reaction either uses or generates an electric
current.
 Voltaic/Galvanic cell: An electrochemical cell in
which a spontaneous reaction generates an electric
current. Reaction occurs by itself.
 Electrolytic cell: An electrochemical cell in which
an electric current is use to drive a chemical
reaction which is not spontaneous. Reaction does
not occur on own.
4
Oxidation-Reduction Reactions
• A cell physically separates an oxidation-reduction
reaction into two half-reactions.
• The force with which electrons travel from the
oxidation half-reaction to the reduction half-reaction is
measured as voltage (actually produces electricity).
red : ( Ag
Ag(s) Cu
AgNO3(aq)
e-

( aq )

 1e  Ag ( s ) )2
2

ox
:
Cu

Cu

2
e
( aq )
(s)
________________________
Cu2+
Cu  Cu2+
Ag+ Ag
ox
red
Cu ( s )  2 Ag  ( aq )  Cu 2 ( aq )  2 Ag ( s )
electrochemical cell
5
CHEMICAL CHANGE  ELECTRIC
CURRENT
Zn metal
Cu2+ ions
With time, Cu plates out
onto Zn metal strip, and
Zn strip “disappears.”
 Zn is oxidized and is the reducing agent
Zn(s)  Zn2+(aq) + 2e Cu2+ is reduced and is the oxidizing agent
Cu2+(aq) + 2e-  Cu(s)
6
Voltaic or Galvanic Cells
anode
oxidation
cathode
reduction
spontaneous
redox reaction
19.2
Chemical to Electrical Energy: Voltaic Cells
• Voltaic cell consists of two half-cells that are
electrically connected.
• Half-cell: A portion of the electrochemical cell where
a half-reaction takes place.
• A simple half-cell can be made from a metal strip
dipped into a solution of its metal ion.
• For example, the silver-silver ion half cell consists of
a silver strip dipped into a solution of a silver salt.
electrode
Cu
Ag
AgNO3
8
Voltaic cell consisting of Cadmium and
Silver electrodes
volt meter
or light
reduction
oxidation
KCl
Cd(NO3)2
AgNO3
Ebbing, D. D.; Gammon, S. D. General Chemistry, 8th ed., Houghton Mifflin, New York, NY, 2005.
9
Voltaic / Galvanic Cells
• The two half-cells are connected to allow electrons to flow
through the external circuit (Metal to Metal Electrode).
• Ag is reduced and Cd is oxidized.
Cathode red: [Ag+ + 1e-  Ag (s)] 2 (must happen twice)
Anode ox: Cd (s)  Cd2+ + 2 eOverall: 2Ag+ + Cd (s)  2 Ag (s) + Cd2+
• The net reaction occuring in the voltaic cell is called the cell
reaction.
• Note that electrons are generated at the anode (-) by
oxidation and thus flow from it to the cathode (+) where
reduction occurs.
10
Voltaic / Galvanic Cells
• Electrons flow from one electrode to the other as long as
there is an external circuit.
• Cadmium has a greater tendency to lose electrons than
silver so atoms in the cadmium electrode lose electrons to
form cadmium ions.
• The electrons flow through the external circuit to the
silver electrode where silver ions gain the electrons to
become silver metal.
• The anode (oxidation) in a voltaic cell has a negative sign
because electrons flow from it.
• The cathode (reduction) in a voltaic cell has a positive
sign
11
Voltaic / Galvanic Cells
• Ions need to flow complete the circuit. This is done
by connecting the two half-cells internally (using a
salt bridge).
• In the absence of the internal connection (salt
bridge), too much positive charge builds up in the
cadmium half-cell (and too much negative charge in
the silver half-cell) and the reaction stops.
• The excess charges at the electrodes must be
balanced by counter ions. It is for this reason that the
salt bridge is used.
• Salt bridges in most cases are made up of KCl.
12
wire
ANODE
OXIDATION
CATHODE
REDUCTION
elect rons
Zn
Zn2+ ions
salt
bridge
Cu
Cu2+ ions
• Electrons travel through the external wire.
• Salt bridge allows anions and cations to move
between electrode compartments.
• This maintains electrical neutrality. A balance
between the positive and negative charges.
13
Two electrodes are connected by
an external circuit.
ox
Cl- Cl-
red
K+ K+
NO3NO3Cu(NO3)2
Ebbing, D. D.; Gammon, S. D. General Chemistry, 8th ed., Houghton Mifflin, New York, NY, 2005.
14
Voltaic Cells
• A salt bridge is a tube of an electrolyte in a gel
that is connected to the two half-cells of a cell.
• The salt bridge allows the flow of ions but
prevents the mixing of the different solutions
that would allow direct reaction of the cell
reactants.
• completes circuit
• keeps electrostatic neutral
15
Voltaic Cells
• The difference in electrical potential between the anode
and cathode is called:
• Cell voltage
• Electromotive force (emf)
• Cell potential
• The potential difference measures the tendency of the
cell reaction to proceed toward equilibrium. It drives
the reaction towards equilibrium.
• Potential difference decreases until it approaches zero
as equilibrium is approached.
• Therefore the amount of electricity produced is finite
(stops at equilibrium).
16
Electrochemical Cell Notation
 It is convenient to have a shorthand way of
designating particular cells.
 Consider the representation of the cell below
2

Cd ( s ) | Cd (aq) || Ag (aq) | Ag ( s )
anode
(oxidation)
salt bridge
cathode
(reduction)
 The anode (oxidation half-cell) is written on the left.
The cathode (reduction half-cell) is written on the right.
 Double line indicates salt bridge is present but not
always present; both half reactions can be in same
container.
17
Electrochemical Cell Notation
2
2
Zn(s ) | Zn (aq ) || Cu (aq ) | Cu(s )
salt bridge
anode
cathode
• The cell terminals are at the extreme ends in the cell
notation (electrode metal).
• A single vertical bar indicates a phase boundary, such as
between a solid terminal and the electrode solution (states
usually omitted).
red : Cu
2
( aq )

 2e  Cu ( s )
ox : Zn( s )  Zn
2
( aq )
 2e

Zn( s )  Cu 2 ( aq )  Cu ( s )  Zn 2 ( aq )
18
Electrochemical Cell Notation
 When the half-reaction involves a gas, an inert
material such as platinum or carbon serves as a
terminal and the electrode surface on which the
reaction occurs.
 Example is the hydrogen electrode; hydrogen bubbles
over a platinum plate immersed in an acidic solution.
 The cathode half-reaction is therefore


2H (aq )  2e  H 2 ( g )
19
Electrochemical Cell Notation
The notation for the hydrogen electrode, written as
a cathode, is
anode

||
H (aq ) | H 2 (g ) | Pt
Sometimes written as
SHE (1atm, 1M)
To write such an electrode as an anode, you simply
reverse the notation; want terminal as extreme end.

cathode
Pt | H 2 (g ) | H (aq ) ||
20
Electrochemical Cell Notation
 To fully specify a cell, it is necessary to give the
concentrations of solutions and the pressure of gases.
 In the cell notation, these are written in parentheses. For
example,
2

Zn | Zn (1.0 M ) || H (1.0 M ) | H 2 (1.0 atm) | Pt
eZn
+
ox
anode
Zn  Zn2+ +2e-
Eo
red:
2H+ +2e-  H2
Eo
Pt
H2 1atm
Zn2+ 1M
ox:
H+ 1M
red
cathode
2H+ + Zn  Zn2+ + H2 Eocell
Note: spectator ions are not usually present in short notation.
This is an example of a standard cell: 1M, 1atm, 25oC (298K)
21
Electromotive Force
• Potential difference Ecell: The difference in
electric potential (electrical pressure) between
two points.
• It is measured in volts.
• The volt, V, is the SI unit of potential difference
equivalent to 1 joule of energy per coulomb of
charge.
1 volt  1 J
C
22
Electromotive Force
• The Faraday constant, F, is the magnitude of
charge on one mole of electrons; it equals 96,500
coulombs (9.65 x 104 C).
1 F = 96,500 C = charge of 1 mole e• In moving 1 mole of electrons through a circuit,
the numerical value of the work done by a cell is
the product of the Faraday constant (F) times the
potential difference between the electrodes.
work(J)   F(coulombs )  volts(J/coulomb )
work done by the system
23
Electromotive Force
• In the normal operation of a voltaic cell, the potential
difference (voltage) across the electrodes is less than
the maximum possible voltage of the cell.
• The actual flow of electrons reduces the electrical
pressure; therefore, slightly lower voltage obtained.
• Electromotive force (EMF) of the cell, or Ecell is the
maximum potential difference between the electrodes
of a cell.
• It can be measured by an electronic digital voltmeter.
• The electrode potential is an intensive property
whose value is independent of the amount of
species in the reaction.
24
Standard Cell EMF’s and 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, while the reaction at the cathode has a
definite reduction potential.
 The overall cell EMF is a combination of these two
potentials.
Ecell = reduction potential, Ered + oxidation potential, Eox
25
Standard Cell emf’s and Standard Electrode
Potentials
Reduction potential, Ered:
• A measure of the tendency to gain electrons in the
reduction half-reaction.
Ered A+ + 1e-  A (s)
E=xV
Oxidation potential, Eox:
• A measure of the tendency to lose electrons in the
oxidation half-reaction.
• For an oxidation half-reaction it is the negative of
the reduction potential for the reverse reaction.
Eox A (s)  A+ + 1e- E = -x V
26
Standard Electrode Potentials
 By convention, the Table of Standard (1M, 1atm,
25oC) Electrode Potentials are tabulated as reduction
potentials (all).
 Consider the zinc-copper cell, calculate the Eocell.
2
2
Zn( s ) | Zn (1M ) || Cu (1M ) | Cu ( s )
2
ox : Zn( s )  Zn (aq )  2e
2
 Eo = 0.76V (changed sign from table)

red : Cu (aq )  2e  Cu ( s )
Zn + Cu2+ --> Zn2+ + Cu
Eo = 0.34V
Eocell = 1.10 V
This voltage for standard cell: 1M otherwise need to correct
from std conc of 1M. Note table is electrode potentials for
standard (superscript o: 1M, 1atm, 298K)
27
•
E0 is for the reaction as
written
•
The more positive E0 the
greater the tendency for
the substance to be
reduced
•
The half-cell reactions are
reversible
•
The sign of E0 changes
when the reaction is
reversed
•
Changing the
stoichiometric coefficients
of a half-cell reaction does
not change the value of E0
reducing agent
oxidized
reduced
All are reductions;
Flip and change sign
for oxidation
Standard Conditions:
1M, 1atm, 25oC
H: reference electrode
Ag+  Ag
is spontaneous
oxidizing
agent
Ebbing, D. D.; Gammon, S. D. General Chemistr
29
8th ed., Houghton Mifflin, New York, NY, 2005.
Standard Cell EMF’s and Standard Electrode
Potentials
• Thus, the electrode potential for the half-reaction
below is Eo = 0.34V.
Cu
2

(aq )  2e  Cu(s )
• The reaction below also has the same electrode
potential ie Eo = 0.34V
2Cu
2

(aq )  4e  2Cu(s )
30
Standard Cell Electrode Potentials
 Standard EMF (theoretical potential), E°cell:
This is the EMF of a cell operating under standard
conditions of concentration (1 M), pressure (1
atm), and temperature (25°C).
 Note that individual electrode potentials require
that we choose a reference electrode.
 Arbitrarily assign this reference electrode a
potential of zero and obtain the potentials of the
other electrodes by measuring the EMF’s. These
are relative values not absolute.
31
Tabulating Standard Electrode Potentials
• By convention, the reference chosen for comparing
electrode potentials is the standard hydrogen
electrode (SHE).
• Standard electrode potentials are measured relative
to this hydrogen reference as the anode.
32
Standard Reduction Potentials
• Standard reduction potential (E0) is the voltage
associated with a reduction reaction at an electrode
when all solutes are 1 M and all gases are at 1 atm.
Reduction Reaction
2e- + 2H+ (1 M)
H2 (1 atm)
E0 = 0 V
Standard hydrogen electrode (SHE)
19.3
Tabulating Standard Electrode Potentials
• For example, when measuring the EMF of a cell composed
of a zinc electrode connected to a hydrogen electrode,
you obtain 0.76 V with hydrogen being cathode and zinc as
anode for the spontaneous reaction.
• By definition, we want the comparison of hydrogen to
other species to be hydrogen as anode and the other species
as cathode; Since hydrogen is given zero value all voltage
measured is given to zinc but we must change sign for
reduction value of zinc.
• Since zinc acts as the anode (oxidation) in this spontaneous
cell, its reduction potential is listed as –0.76 V. This means
that zinc has a potential that is 0.76 less than hydrogen
(relative not absolute)
34
Standard Reduction Potentials
E0 = 0.76 V
cell
Standard EMF(E0cell )
0
0
E0 = Ecathode
- Eanode
cell
Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)
𝑜
𝐸𝑐𝑒𝑙𝑙
=
𝑜
𝐸𝐻 + /𝐻2
𝑜
0.76 𝑉 = 0 − 𝐸𝑍𝑛
+ /𝑍𝑛
Zn2+ (1 M) + 2e-
−
𝑜
𝐸𝑍𝑛+ /𝑍𝑛
𝑜
𝐸𝑍𝑛
+ /𝑍𝑛
Zn
= −0.76 𝑉
E0 = -0.76 V
19.3
Predicting the Spontaneous Direction
of a Redox Reaction
• Standard electrode potentials are useful in determining the
strengths of oxidizing and reducing agents under standardstate conditions.
• The strongest oxidizing agents (species undergoes
reduction) in a table of standard electrode potentials are the
species corresponding to the half-reactions with the largest
(most positive) Eo values. (For example F2(g) highest E in
the table and strongest oxidizing agent in the table -- prefers
to be reduced compared to any other species in table. Will
always be reduced in spontaneous reaction with any of the
species in table.)
• Bottom line: Larger Eo, stronger oxidizing agent, more
tendency to undergo reduction with other species.
36
Strengths of Oxidizing and Reducing
Agents
• Consequently, the strongest reducing agents (species
undergoes oxidation) in a table of standard electrode
potentials are the species corresponding to the halfreactions with the smallest (most negative) Eo
values. (for example, Li smallest E in the table;
therefore, Li (s) strongest reducing agent in the table
[note: must flip for oxidation] -- prefers to be oxidized
compared to any other species in table. Will always be
oxidized in spontaneous reaction with any of species in
table.)
• Bottom line: Smaller Eo, stronger reducing agent, more
tendency to undergo oxidation with other species.
37
Calculating Cell EMF’s from Standard
Potentials
What would be the spontaneous reaction between Cd and
Ag? Calculate Eocell for the spontaneous reaction at 25oC
and 1 M (standard cell)
– Consider a cell constructed of the following two
half-reactions (given from table)
2

o

Reducing agent
Cd (aq )  2e  Cd(s ); E  0.40 V

Ag (aq )  1e  Ag(s ); E  0.80 V
o
Oxidizing agent
Ag+ higher E therefore expect to be reduced (oxidizing agent)
Cd (s) oxidized (reducing agent) in spontaneous reaction.
38
Calculating Cell EMF’s from Standard
Potentials
• Therefore, you reverse the half-reaction and change the
sign of the half-cell potential of cadmium.
2

ox : Cd ( s )  Cd (aq )  2e ; E  0.40 V


o
red : Ag (aq )  1e  Ag ( s ); E  0.80 V
o
• We must double the silver half-reaction so that when the
reactions are added, the electrons cancel.
• This does not affect the half-cell potentials, which do
not depend on the amount of substance.
39
Calculating Cell EMF’s from Standard
Potentials
• Now we can add the two half-reactions to obtain
the overall cell reaction and cell EMF.
• Note: positive voltage meaning spontaneous
reaction for standard cell. If not standard cell
must correct and will discuss later.
2

ox : Cd ( s )  Cd (aq )  2e ;


red : 2 Ag (aq )  2e  2 Ag ( s );
E  0.40 V
o
E  0.80 V
o
Cd(s )  2 Ag  (aq )  Cd 2 (aq )  2 Ag(s ); Eocell  1.20 V
40
Calculating Cell emf’s from
Standard Potentials
• How would we write and draw the cell we just did
in short notation?
2

Cd(s ) | Cd (1M ) || Ag (1M ) | Ag(s )
e-
+
Cd
Cd2+ 1M
ox
anode
Ag
Ag+ 1M
red
cathode
41
A Problem To Consider
 Calculate the standard emf, Eocell, for the
following cell at 25°C.
3
2
Al ( s ) | Al (1M ) || Fe (1M ) | Fe( s )
ox
red
– The reduction half-reactions and standard
potentials are (given)
3

o
2

o
Al (aq )  3e  Al(s ); E  1.66 V
Fe (aq )  2e  Fe(s ); E  0.41 V
42
A Problem To Consider
Given
Al 3 (aq )  3e   Al(s ); Eo  1.66 V
Fe 2 (aq )  2e   Fe(s ); Eo  0.41 V
3
2
Al ( s ) | Al (1M ) || Fe (1M ) | Fe( s )
– You reverse the first half-reaction and its half-cell potential
to obtain
ox : Al ( s )  Al 3 (aq )  3e  ; E o  1.66 V
red : Fe 2 (aq )  2e   Fe( s ); E o  0.41 V
43
A Problem To Consider
3
2
Al ( s ) | Al (1M ) || Fe (1M ) | Fe( s )
2X
3X
3

ox : Al ( s )  Al (aq )  3e ; E  1.66 V
2
o

red : Fe (aq )  2e  Fe( s ); E  0.41 V
o
– To obtain the overall reaction we must balance the
electrons.
3

2 Al(s )  2 Al (aq )  6e ; E  1.66 V
o
3Fe 2 (aq )  6e   3Fe(s ); Eo  0.41 V
note: no multiply by factor to E
44
A Problem To Consider
• Now we add the reactions to get the overall cell
reaction and cell EMF.
3

2 Al(s )  2 Al (aq )  6e ; E  1.66 V
2

o
3Fe (aq )  6e  3Fe(s ); E  0.41 V
o
2Al(s )  3Fe 2 (aq)  2Al 3 (aq)  3Fe(s ); Eo  1.25 V
reducing
agent
oxidizing
agent
• Spontaneous implies +V obtained for standard cell
45
Lecture Presentation
CHEM 111
Thank
You