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Introduction to electrochemistry by t. hara
1. Introduction to Electrochemistry
1. Definition of electrochemistry
1.1. Conversion of chemical energy into electrical energy
1.2. Conversion of electrical energy into chemical energy
1.3. Secondary battery (Rechargeable battery)
1.4. Corrosion
1.5. Conversion of photon energy into electrical energy
via electrochemical reaction
1
Electrical
Energy
Chemical
Energy
2. Introduction to Electrochemistry
1. Definition of Electrochemistry
Electrochemistry deals with
- the conversion of chemical energy
into electrical energy.
- the conversion of electrical energy
into chemical energy.
2
Electrical
Energy
Chemical
Energy
4. Introduction to Electrochemistry
4
Chemical
Energy?
* Fermi level, EF
[eV]=[1.6*10-19C·V]=[1.6*10-19C·J/C]
=[1.6*10-19J]
* Standard reduction potential
[V] =[J/C]
Electrical
Energy?
Energy [J]
= Potential [V=J/C]*Charge [C]
= Potential [V=J/C]*Current [A=C/s]
*Time [s]
5. Introduction to Electrochemistry
Note:
Fermi level, EF, is equal to internal chemical potential at 0
K, μ0. But in electrochemistry, we have to take some effects
into consideration: solvation, counter ion, viscosity of
solution etc. can be listed as possible causes that can have
an influence on the chemical potential. In electrochemistry,
established theories such as Debye-Huckel limiting law
(logγ+-=-Bz+|z-|I1/2, γ is the activity coefficient, B is the
constant, z is the valency of ion, I is the ionic strength that
is caluculated as I=(1/2)Σcizi
2).
Works only for dilute solutions (typically <= 0.01 M) of
strong electrolytes that can be completely ionized.
6. Introduction to Electrochemistry
6
Electrochemical reaction
= Redox reaction
= Reduction and Oxidation
= Reactions in which transfers of electrons take place.
Reduction
= the process in which electrons are gained by a reactant.
Oxidation
= the process in which electrons are donated by a reactant.
7. Introduction to Electrochemistry
1.1. Conversion of chemical energy
into electrical energy.
One of the applications is primary battery (Galvanic cell).
Galvanic cell is one in which electrical energy is
spontaneously produced by chemical reactions.
7
Electrical
Energy
Chemical
Energy
8. Introduction to Electrochemistry
This is a cartoon of
a primary cell, which is
called Daniell cell, which
was invented in the
1830’s by the British
chemist Daniell.
A zinc bar is placed
into a ZnSO4 solution,
a copper bar is placed
into a CuSO4 solution.
8
Zn
Cu
wire
Salt bridge
Cu(II) sulfateZn(II) sulfate
electrons
9. Introduction to Electrochemistry
The zinc bar and the copper bar are connected with
conducting wire. Free mobile electrons flow from the
anode to cathode through the wire (and external
circuits, such as a flush light.).
9
The zinc sulfate solution
and the copper sulfate
solution are connected via
salt bridge. Ions in the
electrolyte transfer and
balance the charge via salt
bridge.
Zn
Cu
wire
Salt bridge
Cu(II) sulfateZn(II) sulfate
electrons
10. Introduction to Electrochemistry
Because of the potential difference
between zinc and copper,
electrons are going to flow spontaneously
through the conducting wire,
resulting in the oxidation of zinc metal to zinc cation
and the reduction of copper cation to copper metal.
What does it mean?
10
11. Introduction to Electrochemistry
Potential ? Potential difference?
… Thermodynamic driving forth!
11
A battery can have electromotive force (emf), the same
as electric potential difference between two
electrodes.
If two electrodes have the different potentials, the
battery can give us electricity until the two electrodes
have the same potential = Gibbs free energy difference
goes to zero = the Fermi levels are aligned.
12. Introduction to Electrochemistry
12
An equilibrium electrochemical potential is described as
a standard reduction potential.
Zn2+(aq) + 2e- → Zn(s) -0.76 V
Cu2+(aq) + 2e- → Cu(s) +0.34 V
13. Introduction to Electrochemistry
13
Zn2+(aq) + 2e- → Zn(s) -0.76 V
Standard reduction potential?
A measure of the tendency of a chemical species to
acquire electrons.
Cu2+(aq) + 2e- → Cu(s) +0.34 V
The more negative the potential, the greater the
species' ability to donate electrons and tendency to be
oxidized.
14. Introduction to Electrochemistry
14
Cu2+(aq) + 2e- → Cu(s) +0.34 V
Zn2+(aq) + 2e- → Zn(s) -0.76 V
SRP
E=1.10 V
Thanks to the potential difference, you will gain an
emf, E (or total cell potential) of 1.10 V.
It means this reaction occurs spontaneously.
15. Introduction to Electrochemistry
15
+0.34 V
-0.76 V
SRP
E=1.10 V
ΔG=-212 kJ/mole
Energy has a unit, Joule [J]
[J]=[V·C]=[V·A·s]
ΔG=-nFE (F=96485 C/mole)
[J/mole]=[C/mole][V]
n is the number of electrons
involved in the reaction.
Negative free energy change
(ΔG<0 because E>0)
is identified as defining a
spontaneous process.
16. Introduction to Electrochemistry
16
Note:
The output potential (potential
difference between an anode and
cathode) of a primary cell gives the
maximum potential at zero current
flow.
Once electrons are allowed to flow
through the circuit, the actual output
potential changes with time, because
the driving force for the reaction
decreases as the system approaches
equilibrium.
Secondary battery can be recharged
and regain its electromotive force.
Potential
More negative potential:
It push electrons
through the external
circuit.
17. Introduction to Electrochemistry
17
Note:
- All thermodynamic measurements are of differences
between states. There is no absolute value for any
thermodynamic property, except for entropy.
In order to quantify thermodynamic values (in
electrochemistry),
(1) a temperature is chosen at 298 K (25 °C),
(2) a pressure is chosen at 1 atm (105 Pa),
(3) a concentration is chosen at 1 mol/L.
These conditions are called standard conditions.
18. Introduction to Electrochemistry
18
When cell is not at standard conditions, use Nernst
equation:
E=E0-(RT/nF)lnK
where, E0 is the total cell potential. So far, we have
omitted subscript zero, because we have taken standard
conditions for granted. R is gas constant 8.315 J/Kmol, T
is temperatures in Kelvins, and K is reaction quotient. In
Daniell cell, K=[Zn2+]/[Cu2+]. Note that K is the ratio of
[product]the number of moles of the product to [reactant]the number of
moles of the reactant. As E declines with reactants converting
products, E eventually reaches zero. Zero potential means
reactions is at equlibrium, namely, battery is dead.
19. Introduction to Electrochemistry
19
Reduction:
Cu2+(aq) + 2e- → Cu(s)
Oxidation:
Zn(s) → Zn2+(aq) + 2e-
Total Cell Reaction:
Zn(s)+Cu2+(aq) → Zn2+(aq)+Cu(s)
Electrochemical reaction is composed of two half
reactions, namely, oxidation and reduction reactions.
Copper cation is
being reduced
(its oxidation
number is going
from +2 to 0).
Zinc is being
oxidized (its
oxidation number
is going from 0 to
+2).
Zn
Cu
wire
Salt bridge
Cu(II) sulfateZn(II) sulfate
electrons
20. Introduction to Electrochemistry
20
Oxidation number represents the number of electrons
required to produce the charge on a species: (1) to be
oxidized is to lose electrons (e.g., Zn(s) → Zn2+(aq) + 2e-);
(2) to be reduced means to gain electrons (Cu2+(aq) + 2e-
→ Cu(s)). Here, (s) stands for the solid state, (aq) stands
for aqueous ion.
The oxidation number...
(1) for any elemental substance is zero.
(2) for an ion is its charge (e.g., Zn2+ has +2).
21. Introduction to Electrochemistry
In electrochemical terminology, an electrode at which an
oxidation reaction occurs is called an anode. An electrode
at which a reduction reaction occurs is called a cathode.
21
Cathode Reaction:
Cu2+(aq) + 2e- → Cu(s)
Anode reaction:
Zn(s) → Zn2+(aq) + 2e-
Oxidation Reduction
Zn
Cu
wire
Salt bridge
Cu(II) sulfateZn(II) sulfate
electrons
22. Introduction to Electrochemistry
Zn/Zn2+ pair has a more negative standard reduction
potential than Cu/Cu2+ has (the former is going to be
oxidized, the latter is going to be reduced.).
22
Cathode Reaction:
Cu2+(aq) + 2e- → Cu(s)
Anode reaction:
Zn(s) → Zn2+(aq) + 2e-
Cu2+(aq) + 2e- → Cu(s) +0.34 V
Zn2+(aq) + 2e- → Zn(s) -0.76 V
SRP
E=1.10 V
23. Introduction to Electrochemistry
23
+0.34 V 4.78 eV
-0.76 V 3.68 eV
SRP EF
We can also use Fermi level,
that is used in solid state physics
and semiconductor physics.
After understanding the
relationship between the
different branches of science,
you will be able to use more
resources of knowledge.
24. Introduction to Electrochemistry
24
Zn2+(aq) + 2e- → Zn(s) 3.68 eV
Fermi level?
A minimum energy to remove electron from a material.
Cu2+(aq) + 2e- → Cu(s) 4.78 eV
The more positive the level, the greater the species'
ability to donate electrons and tendency to be
oxidized.
26. Introduction to Electrochemistry
26
SRP
0 eV-4.44 V
0.00 V 4.44 eV
Standard Hydrogen Electrode
2H++2e-->H2
+1.23 V 5.63 eV
Standard Oxygen Electrode
O2(g) + 4H+(aq) + 4e- → 2H2O(l)
In electrochemistry, the
standard hydrogen
electrode (SHE) potential is
taken as a reference point.
EF
27. Introduction to Electrochemistry
27
SRP
0 eV-4.44 V
0.00 V 4.44 eV
+1.23 V 5.63 eV
This is the vacuum level
(Evac) at which electrons
are at rest in vacuo, just
outside the surface of the
electrode.
It is taken as a reference
point in solid state
physics.
The closer an electron is
to the vacuum level, the
weaker it is bound to the
solid.
EF
28. Introduction to Electrochemistry
28
In electrochemical reactions, electrons just change
Places. Charge is conserved.
In addition, a properly balanced redox reaction means
both mass and charge are conserved.
Cathode Reaction:
Cu2+(aq) + 2e- → Cu(s)
Anode reaction:
Zn(s) → Zn2+(aq) + 2e-
Total Cell Reaction:
Zn(s)+Cu2+(aq) → Zn2+(aq)+Cu(s)
29. Introduction to Electrochemistry
29
During discharging a primary battery, free mobile
electrons flow from the anode to cathode through the
wire, ions in the electrolyte transfer and balance the
charge.
Zn
Cu
wire
Salt bridge
Cu(II) sulfateZn(II) sulfate
electrons
32. Introduction to Electrochemistry
32
Total Cell Reaction:
Zn(s)+Cu2+(aq) → Zn2+(aq)+Cu(s)
Zn2+(aq)
increased.
e-
Zn(s)
e-
Cu(s)
Cathode Reaction:
Cu2+(aq) + 2e- → Cu(s)
Anode reaction:
Zn(s) → Zn2+(aq) + 2e-
Salt bridge. K+(aq) etc.
33. Introduction to Electrochemistry
33
Salt bridge?
A tube or membrane packed
with a solution of salt
composed of ions not involved
in the cell reaction. The ions just
permit exchange of charge in
order to balance the charge.
Zn
Cu
wire
Salt bridge
Cu(II) sulfateZn(II) sulfate
electrons
34. Introduction to Electrochemistry
34
Shorthand notation of cell.
cathode | catholite || anolite | anode
Zn
Cu
wire
Salt bridge
Cu(II) sulfateZn(II) sulfate
electrons
Shorthand notation of Daniell cell.
Zn(s) | Zn2+(aq, 1M) ||Cu2+(aq, 1M)| Cu(s)
35. Introduction to Electrochemistry
Note:
Assume we have m g (N mole) of Zn with A.M.=65 g/mol,
the required I·t (t is the duration time [s] ) is calculated as
follows:
N=m/A.M.,
N=s·I·t/(n·F),
I·t=N·n·F/s=(m/65)·2·F/1,
where s is the stoichiometric coefficient of the species.
In the following reaction, the s of Zn(s) is one.
Zn(s)+Cu2+(aq) → Zn2+(aq)+Cu(s)
35
36. Introduction to Electrochemistry
36
The LeClanché cell, often called dry cell (but it is not “dry,”
it uses gel electrolyte, and it can leak.), is commercially
available primary battery. The reactions involved are:
anode: Zn (s)->Zn2++2e–
cathode: MnO2(s)+H2O+NH4+e-->Mn(OH)3(s)+NH3(aq)
The output voltage is 1.55 V.
Alkaline battery uses the same reactant as above, but
under basic (alkaline) conditions.
37. Introduction to Electrochemistry
SUMMARY of 1.1.
chemical energy -> electrical energy
We have seen how to construct a primary cell that is
capable of generating a spontaneous flow of electrons.
The flow of electrons (current) can be used to perform
work on electronic appliances.
A spontaneous flow of electrons is induced by
electromotive force (emf), potential difference between
an anode and a cathode.
Recall that Potential [V]=Energy [J]/Charge (C).
37
38. Introduction to Electrochemistry
1.2. Conversion of electrical energy
into chemical energy.
One of the applications is electroplating.
38
Electrical
Energy
Chemical
Energy
39. Introduction to Electrochemistry
This is the world oldest
electroplating equipment
that found in an ancient
tomb in Bagdad in 1936. It
consists of a 14-centimeter-
high egg-shaped clay jar with
an asphalt stopper. An iron
rod protruding out of the
asphalt is the anode, which is
surrounded by a copper
cylinder used as the cathode.
Filled with vinegar as an
electrolytic solution.
39
40. Introduction to Electrochemistry
Electroplating of sacrificial anode (nonspontaneous)
40
Zn
Metal
e.g.,
steel
Zn(II) sulfate
electrons
Zn(s) deposition
Zn(s)
Zn2+(aq)
e-
Spontaneous reaction
Zn(s)+Fe2+(aq) → Zn2+(aq)+Fe(s) E=0.32 V
41. Introduction to Electrochemistry
Control of electrochemical potentials of electrodes allows
the reaction to be controlled (even if the reaction cannot
be occur spontaneously).
41
Potential
Zn
Metal
Zn(II) sulfate
electrons
Zn deposition
0.76 V
>0.76 V
opposite to
spontaneous reaction
42. Introduction to Electrochemistry
42
Energy [J]=[V·C]=[V·A·s]
ΔG [J/mole]=-nFE [C/mole][V]
Positive free energy change (ΔG>0 because E<0)
is identified as defining a nonspontaneous process.
Zn
Metal
Zn(II) sulfate
electrons
Zn(s) deposition
Zn(s)
Zn2+(aq)
e-
43. Introduction to Electrochemistry
Note:
To electroplate m g (N mol) of metal with F.W.=M g/mol,
the required I·t (t is the duration time [s] ) is calculated as
follows:
N=m/M,
N=s·I·t/(n·F),
I·t=N·n·F/s=(m/M)·n·F/s,
where s is the stoichiometric coefficient of the species.
In the following reaction, the s of Zn(s) is one.
Cu(s)+Zn2+(aq)->Cu2+(aq)+Zn(s)
43
44. Introduction to Electrochemistry
44
Primary Battery
Chemical Energy
-> Electrical Energy
ΔG<0 , spontaneous
Electroplating
Electrical Energy
-> Chemical Energy
ΔG>0 , nonspontaneous
SUMMARY of 1.2.
electrical energy -> chemical energy
With an external voltage in the opposite direction of
spontaneous reaction, electrical energy is converted into
chemical energy.
46. Introduction to Electrochemistry
(1) During charging, electrical energy is converted into
chemical energy.
The charging is conducted by applying an external
voltage of opposite polarity to that of discharging in order
to gain higher total cell potential (=higher energy).
(2) During discharging, chemical energy is converted into
electrical energy.
So we can use the stored energy in the similar way as
primary batteries.
46
51. Introduction to Electrochemistry
51
Lithium ion battery
C+LiCoO2 ↔ LiC6+Li0.5CoO2 E=3.6 V
Capacity/cell:
e.g., 2.2 [Ah]=2.2 [C/s]*3600 [s]
=7920 [C]
Energy density/cell:
e.g., 3.6 [V]*2.2 [Ah]=7.92 [Wh]
=3.6 [J/C]*7920 [C]=102643.2 [J]≈103 [kJ]
One Ah is defined as one ampere that is passed for one hour.
52. Introduction to Electrochemistry
1.4. Electrochemical (Galvanic) corrosion
Corrosion is the gradual destruction of materials
(usually metals) by chemical reaction.
52
55. Introduction to Electrochemistry
The metal surfaces (except for Au) are covered with oxide
films in oxidative atmosphere (e.g., air). Some oxide films
are brittle and easily peeled off the metal surfaces. If a
metal with a surface oxide film and a bare metal surface
coexist and they have electrical contact, corrosion occurs.
55
Steel (Fe)
Fe2O3
Seawater
56. Introduction to Electrochemistry
When the oxide-free surface of a metal becomes exposed
to the solution, positively charged metal ions tend to pass
from the metal into the solution, leaving electrons behind
on the metal.
56
Steel (Fe)
Fe2O3
Seawater
Anode reaction:
Fe(s) → Fe2+(aq) + 2e-
Fe2+
e-
57. Introduction to Electrochemistry
The concentration of dissolved oxygen in air saturated
aqueous solutions at ambient temperature is about 2
mM. Oxygen that reaches at the surface of the steel gains
electrons from steel, and is reduced to hydroxide anion.
57
Steel (Fe)
Fe2O3
Seawater
Fe2+
e-
O2
OH-
O2
O2
O2
Cathode Reaction:
O2 + 2H2O + 4e- → 4OH-
O2
58. Introduction to Electrochemistry
In other words. the accumulation of negative charge on
the steel surface due to the residual electrons leads to an
increase in the potential difference between the metal
and the solution. This change in the potential encourage
the deposition of dissolved metal ions from the solution
onto the metal.
58
Steel (Fe)
Fe2O3
Seawater
Fe2+O2
OH-
Fe2++2OH- ->Fe(OH)2
e-
FexOy·nH2O
59. Introduction to Electrochemistry
59
Cathode Reaction:
O2 + 2H2O + 4e- → 4OH-
Anode reaction:
Fe(s) → Fe2+(aq) + 2e-
+0.40 V 4.84 eV
-0.44 V 4.00 eV
SRP
The system has an emf of 0.84 V.
This reaction occurs spontaneously.
EF
60. Introduction to Electrochemistry
The concentration of dissolved oxygen is low and so the
rate of transport of oxygen often limits the cathodic
reduction current and the corrosion rate. Under these
conditions the corrosion rates depend only on the rate of
reduction of the cathodic reactant and the corrosion is
said to be under cathodic control.
Note that in almost all the elctrochemical reaction the
rate determining step is not the electron propagation but
the mass transfer, i.e., molecular or ion diffusion.
60
62. Introduction to Electrochemistry
The recent years dye sensitized solar cells or Grätzel cells
have attracted considerable attention worldwide due to
their mechanism that is different from “conventional”
semiconductor-based (Si, GaAs etc.) solar cells.
62
63. Introduction to Electrochemistry
Energy conversion in a Grätzel cell is based on the
injection of an electron from a photoexcited state of the
sensitizer dye into the conduction band of a
nanocrystalline oxide semiconductor (anatase TiO2 etc.).
63
Measured .
StandardReductionPotential(V)
-0.5
0
0.5
1.0
SRP EF
TiO2 (-0.44 V) 4.0 eV
Dye* ( 0.74 V) 3.7 eV
Dye ( 1.06 V) 5.5 eV
I3
-/I- 0.55 V (3.9 eV)
e-
h
Conducting glass
electrode
64. Introduction to Electrochemistry
The oxidized dye is reduced and regenerated to its
ground state by a liquid electrolyte redox couple (I3
-/I-
etc.). Regeneration of iodide ions to tri-iodide is achieved
at a counter electrode.
64
Measured .
StandardReductionPotential(V)
-0.5
0
0.5
1.0
SRP EF
TiO2 (-0.44 V) 4.0 eV
Dye* ( 0.74 V) 3.7 eV
Dye ( 1.06 V) 5.5 eV
I3
-/I- 0.55 V (3.9 eV)
e-
h
Conducting glass
electrode
65. Introduction to Electrochemistry
Light absorption is accomplished by a monolayer of
photoactive dye adsorbed chemically at the TiO2 surface
and excited by interaction with an incident photon of
light.
65
Conducting glass
electrode
TiO2
h
Notas del editor
Today, I will talk about introductory electrochemistry.
Today’s topic can be shown here.
Starting with the definition of electrochemistry, we will discuss 1.1. Conversion of chemical energy into electrical energy, then, 1.2. Conversion of electrical energy into chemical energy. After this, we will be able to understand cyclic energy conversion: 1.3. Secondary battery (Rechargeable battery). 1.4. Corrosion is also useful to know about electrochemical phenomena in our daily life. In addition, I will introduce 1.5. Conversion of photon energy into electrical energy via electrochemical reaction. This requires the understanding semiconductor physics besides electrochemistry. I would like to emphasize that you all should find and make new area of science or technologies. To understand different branches of science, to combine those stuff, and to restructure the new area will be definitely useful for new generation, although I will be going to be categorized into old generation in the near future or already?
Let me start with the definition of electrochemistry. With defining words, terms, or concepts, we can prepare for upcoming cumbersome and complicated talk.
As mentioned earlier, electrochemistry deals with
- the conversion of chemical energy into electrical energy.
- the conversion of electrical energy into chemical energy.
What is energy?
The useful definition is … to be unstable is to have an ability to work, to influence others, or to change something.
In applied electrochemistry, we often make unstable states intentionally and use it for some purposes.
What is chemical energy?
One useful expression is Fermi level, EF, it is equal to internal chemical potential at 0 K.
Its unit is [eV]=[1.6E-19C·V]=[1.6E-19C·J/C]=[1.6E-19J]
Fermi level is used in semiconductor physics, but to understand its definition will be useful for electrochemist.
In electrochemistry, standard reduction potential is widely used.
Its unit is [V]=[J/C]
Next,
What is electrical energy?
Its unit is [J]=Potential [V=J/C]*Charge [C]=Potential [V=J/C]*Current [A=C/s]*Time [s].
Not to memorize difficult equation. Just use mathematical expression just to pretend to be smart.
But, please memorize above relationships.
Of course, in electrochemistry we will be thinking about electrochemical reaction.
Electrochemical reaction is
Redox reaction,
It means Reduction and Oxidation,
In these reactions, transfers of electrons take place.
Reduction is the reaction in which electrons are gained by a reactant.
Oxidation is the reaction in which electrons are donated by a reactant.
It should be noted that oxidants are going to donate electrons not because they have generous disposition but because they have more negative potentials that push electrons into guys with more positive potentials.
Potential or potential difference between two reactants becomes thermodynamic driving force of electrochemical reactions.
An equilibrium electrochemical potential is described as a standard reduction potential.
Here, we have two reduction reactions. One is the reduction of zinc cation: zinc cation gains two electrons, and then becomes zinc metal, solid zinc. This reduction reaction has a standard reduction potential of minus 0.76 volt. The other is the reduction of copper cation: copper cation gains two electrons, and then becomes copper metal. The standard reduction potential of this reaction is plus 0.34 volt.
Energy has a unit, Joule [J] [J]=[V·C]=[V·A·s]
Gibb’s free energy change which is defined by the equation, ΔG=-nFE [J/mole]=[C/mole][V], shows the thermodynamic energy given off or consumed in the reaction. Here, F is Faraday constant, F=96485 C/mole. One mole of electron carries 96485 coulombs. n is the number of moles of electrons involved in the reaction.
We obtained negative free energy change, because the total cell potential is positive (ΔG<0 because E>0).
Negative free energy change is identified as defining a spontaneous process.
Please don’t leave yourself confused. Positive E, total cell potential, implies spontaneous. In contrast, negative E implies nonspontaneous. E=0 means equilibrium reached, namely, battery is dead. Total cell potential is the difference between two standard reduction potentials. Delta G will be negative in spontaneous reactions.
Electrochemical reaction is composed of two half reactions, namely, oxidation and reduction reactions.
Here we have a total cell reaction: zinc solid plus copper two plus aqueous ion makes zinc two plus aqueous ion plus copper solid.
In the left cell, an oxidation reaction is proceeding: zinc solid is going to be zinc two plus aqueous ion and losing two electrons. Zinc is being oxidized. Its oxidation number is going from zero to plus two.
In the right cell, a reduction reaction is proceeding: copper two plus aqueous ion is gaining two electrons and going to be copper solid. Copper cation is being reduced. Its oxidation number is going from plus two to zero.
Fermi level, EF
[eV]=[1.6*10-19C·V]
=[1.6*10-19C·J/C]
=[1.6*10-19J]
The standard potential of minus 0.76 V means the Fermi level of 3.68 eV.
The standard potential of plus 0.34 V is the Fermi level of 4.78 eV.
In electrochemistry, the standard hydrogen electrode (SHE) potential is taken as a reference point: 0 V.
The half cell reaction of standard hydrogen electrode is 2H++2e-->H2. Here, Fermi level is 4.44 eV.
We can see one more example here, it is the standard oxygen electrode half cell reaction that has a standard reduction potential of plus 1.23 V. Here, its Fermi level is 5.63 eV. Actually, it is not describing the exact definition of Fermi level, because this half cell reaction occurs in aqueous solutions. This means the reaction is under the influence of water molecules.
This is the vacuum level (Evac) at which electrons are at rest in vacuo, just outside the surface of the electrode.
It is taken as a reference point, 0 eV, in solid state physics.
The closer an electron is to the vacuum level, the weaker it is bound to the solid.
We will go back to electrochemical reaction. In electrochemical reactions, electrons just change
Places. Charge is conserved.
Here, there are two half cell reactions: on the left we have an anode reaction, an oxidation reaction; on the right there is a cathode reaction, a reduction reaction. Our anode reaction is Zn(s) → Zn2+(aq) + 2e-. And, our cathode reaction is Cu2+(aq) + 2e- → Cu(s). These two half cell reaction makes a total cell reaction: Zn(s)+Cu2+(aq) → Zn2+(aq)+Cu(s). A properly balanced redox reaction means both mass and charge are conserved.
At the anode surface a zinc atom releases its two electrons and itself is going to dissolve into solution. At the cathode surface a copper cation gains two electrons and deposits onto the cathode surface.
Again, at the anode surface a zinc atom releases its two electrons and itself is going to dissolve into solution. The two electrons change their places from zinc atom to the surface of copper cathode, and a copper cation receives the two electrons, resulting in the copper metal deposition onto the cathode surface. In the left half cell, the number of zinc cation was increased. So the a zinc cation is going to diffuse into this salt bridge and/or a divalent anion or two monovalent anions is or are coming from the salt bridge to the left half cell. On the right half cell, the number of copper cation was decreased. Therefore, an divalent anion or two monovalent anions is or are diffusing to the salt bridge and/or a divalent cation or two monovalent cations is or are going to come from the salt bridge.
Ions can move only in the electrolyte, and electrical charge imbalance are corrected through this salt bridge.
Here, I will show you the shorthand notation of electrochemical cell.
The anode half-cell is written first.
Phase boundary, such as the boundary between solid and liquid, are indicated by a single line.
Salt bridge is indicated by a double line.
1.2. Conversion of electrical energy into chemical energy.
One of the application is electroplating.
Later in this talk, you will be informed about the corrosion (rusting) which is one of the aspect of electrochemistry.
Electroplating is conducted in order to prevent or slow down the rusting of steel. A sacrificial anode made from active metal such as zinc is electroplated onto steel. The zinc is preferentially oxidized in place of the iron. When a sacrificial anode is used, the iron in the steel simply serves as a cathode (it is reduced).
In contrast, when the iron is in contact with some inactive metals such as copper, the iron in the steel serves as an anode, then it corrodes.
The current in a galvanic cell flows in the direction of the spontaneous reaction.
In contrast, in order to make the nonspontaneous reaction occur, we have to apply an external voltage in the opposite direction of spontaneous reaction (more negative than -0.32 V).
Electrical energy is also consumed in order to deposit zinc onto the metal. Recall that Energy [J]=[V·C]=[V·A·s].
Here, I will show how the potentials of two electrodes of a secondary battery change during charging/discharging.
At the charged state, there is a potential difference between two electrodes. So, the battery can spontaneously convert its chemical energy (potential difference) into electrical energy.
After supplying electricity, the battery will become discharged state. At this stage, there is no potential difference. So, the battery cannot supply electricity anymore.
However, we can recharge the battery by applying external voltage. The battery can regain chemical energy (potential difference).
In a practical manner, some batteries are not fully charged and not fully discharged. It is just because the materials used for some batteries, such as lithium ion batteries, are not stable under fully charged and/or fully discharged states. Some crystalline materials can be destructed. Some organic materials can be irreversibly converted into another form of chemicals.
In this slide, I will show you one example of secondary batteries. Lead acid batteries are shown here. Lead acid batteries are currently used for automobile starting, lighting and ignition. During discharging, this half cell reaction occurs at the anode surface, and this half cell reactions occurs at the cathode surface. So, the total cell reaction becomes Pb+PbO2+2H2SO4→2PbSO4+2H2O. At the both electrode surfaces, lead sulfate will be deposited. This reaction will give us an output voltage of 2.1 V. During charging, the reaction occurs in the opposite direction.
Next, I will introduce lithium ion batteries. In a lithium ion battery, lithium ions move from a negative electrode to a positive electrode during discharge and back when charging. Lithium ion batteries use an intercalated lithium compounds as the electrode materials. Lithium ion batteries are common in consumer electronics, such as personal computers, because lithium ion batteries have the highest energy densities among rechargeable batteries, so the size of the battery can be smaller than that of lead acid battery, although the electrochemical reactions in lithium ion batteries are slower than those in lead acid batteries. But the chemical reactions’ rates are enough to supply current to consumer electronics products.
The typical capacity of a lithium ion battery is 2.2 ampere hour. It means the charge of 2.2 coulomb flows every second, and this current can continue to flow for 3600 seconds. So, one lithium ion battery can supply 7920 Coulomb.
The energy density of a cell is equal to 3.6 [V]*2.2 [Ah]=7.92 [Wh]. This is equal to 3.6 [J/C]*7920 [C]. This corresponds to 103 [kJ].
1.4. Electrochemical (Galvanic) corrosion
Corrosion is the gradual destruction of materials
(usually metals) by chemical reaction.
The term corrosion is sometimes also applied to the degradation of plastics, concrete and wood, but generally refers to metals.
Corrosion can be the negative aspect of spontaneous electrochemical reaction.
The most widely used metal is iron (usually as steel).
Since the most corrosion products have little mechanical strength, a severely corroded pieces of metal is useless for its original purpose
The consequences of corrosion are many and varied and the effects of these on the safe, reliable and efficient operation of equipment or structures are often more serious than the simple loss of a material.
For example, reduction of steel thickness resulting in loss of mechanical strength and structural breakdown. When the steel is lost in localized zone, crack can develop.
Corrosion of pipes allows escape of their contents and possible harm to the surroundings. The value of your cars is reduced due to deterioration of appearance.
Electrochemical corrosion occurs between two “electrodes” which have electrical contact with each other and are immersed in a common electrolyte.
However, how can the two zones of the same steel be two electrodes which should have the different potential to allow electrochemical reactions.
Localized corrosion can occur on surfaces where the metal is in a varying condition of stress, access of air, temperature, and so on.
In other words. the accumulation of negative charge on the steel surface due to the residual electrons leads to an increase in the potential difference between the metal and the solution. This change in the potential encourage the deposition of dissolved metal ions from the solution onto the metal.
In the pH range 6.5 – 8.5 corrosion is usually accompanied by the formation of solid corrosion debris from the reaction between the anodic and cathodic products.
Fe2++2OH- ->Fe(OH)2, iron hydroxide
Pure Fe(OH)2 is white but the material initially produced by corrosion is greenish white because of partial oxidation.
Fe(OH)2+H2O+1/2O2->2Fe(OH)3 , hydrated iron oxide
Further hydration and oxidation reactions can occur and reddish rust that eventually forms is a complex mixture. Because the rust is precipitated as a result of secondary reactions it is porous and absorbent and encourage further corrosion.
Here, an anode reaction: iron solid goes to iron aqueous iron cation and donate two electrons to its counterpart. This reaction has the standard reduction potential of minus 0.44 V (the Fermi level of 4.00 eV). Recall that the actual reaction is oxidation but we have to estimate the total cell potential difference from two standard reduction potentials of the reaction which will be an anode reaction and of a cathode reaction. Another one is the cathode reaction. That reaction has the standard reduction potential of plus 0.40 V (the Fermi level of 4.84 eV). The total cell potential difference will be 0.84 eV.