ELECTROCHEMISTRY - ELECTRICAL DOUBLE LAYER

Dr.P.GOVINDARAJ
Associate Professor & Head , Department of Chemistry
SAIVA BHANU KSHATRIYA COLLEGE
ARUPPUKOTTAI - 626101
Virudhunagar District, Tamil Nadu, India
ELECTROCHEMISTRY – ELECTRICAL DOUBLE LAYER

ELECTRICAL DOUBLE LAYER (EDL)
Definition:
• Electrical double layer is a structure in the form of fixed layer and diffused layer
that appears on the surface of an object like electrode or colloidal particle when introduced
in an electrolytic solution

• The surface charge of an object is either positive or negative
• The fixed layer consists of ions adsorbed onto the object due to chemical interactions
• The diffused layer is composed of ions attracted to the surface charge via the columbic
force and electrically screening the first layer i.e., the diffused layer is loosely
associated with the object
• The potential difference between the fixed layer and diffused layer is called zeta
potential (or) electro kinetic potential

Evidence for electrical double layer :
• The stability of colloidal particle is due to the existence of electrical double layer around
the surface of colloidal particles that prevent co-aggulation of colloids
Example:
• Stability of AgI colloidal particles is due to the presence of electrical double layer which
prevent the aggregation of colloidal particles
• Similarly,
The aggregation of fat particles in milk
and hemoglobin molecules in blood is
prevented due to the existence of electrical
double layer

Electro capillary phenomena
• Electro capillary phenomena are the phenomena related to changes in the surface
energy (interfacial tension) of the dropping mercury electrode as the electrode potential
changes (or) the electrolytic solution composition and concentration change. The term
electro capillary is used to describe the change in mercury electrode potential as a function
of the change in the surface or interfacial tension of the mercury determined by the
capillary rise method
𝜕
𝜕𝑉 T= - qM ------- Lippmann equation
i.e., change of interfacial tension with potential difference is the measure of the
charge density on the electrode

Electro kinetic Phenomena:
• It is the phenomena, that occur in colloids, in which movement of one phase with respect
to another phase under the influence of electric field
Types of electro kinetic phenomena
i. Electrophoresis
ii. Electro-osmosis
iii. Streaming potential
iv. Sedimentation potential

ELECTROPHORESIS
Definition:
• Electrophoresis is the motion of charged particle through a liquid medium under the
influence of an applied electric potential
• If an electric potential is applied to a colloid, the charged colloidal particles move
towards the oppositely charged electrode

ELECTROPHORESIS
• Rate determining potential for electrophoretic
movement of particles, is zeta potential
• Shear plane is located at the junction (bb’) of
tightly bounded layer (a-b) and diffusion
layer (b-c)
• Zeta – potential is potential difference between
shear plane and electro neutral region

STREAMING POTENTIAL
• When the liquid is forced to flow through the capillary or the plug, by an applied
pressure, an EMF is developed which is called the streaming potential
• If E is the potential difference developed across a capillary of radius, a and length l, for an
applied pressure difference p, we can write the following

ELECTRO-OSMOSIS
• The movement of the dispersion medium under the
influence of applied potential is known
as electro-osmosis
• The dispersed phase is kept, stationary , the medium is
actually found to move to the
electrode of opposite sign that its own
• When the applied pressure exceeds the zeta potential, that
diffuse layer moves and causes
electro-osmosis

SEDIMENTATION POTENTIAL
• When sedimentation of suspend colloids occurs,
this rise a potential difference
• this technique is the least commonly used for the
determination of zeta potential, because of
several limitations associated with the
measurement and calculation of the zeta
potential

ZETA POTENTIAL
Definition:
• Zeta potential is the electrical potential at the slipping plane. This plane is the interface
which separates mobile fluid from fluid that remains attached to the surface. Zeta potential
is a scientific term for electrokinetic potential in colloidal dispersions.

ZETA POTENTIAL
Measurements of zeta potential:
• The zeta potential of a particle is calculated from electro kinetic phenomena such as
Electrophoresis
Streaming potential
Electro-osmosis
Sedimentation potential

ZETA POTENTIAL
Measurements of zeta potential by Electrophoresis method:
• The movement of sol particles under an applied electric
potential is called electrophoresis or cataphoresis
• The zeta potential is calculated by using the Smoluchowski
equation.
• For spherical particle which can be treated as point charges
as follows

ZETA POTENTIAL
Applications of zeta potential:
• Prepare colloidal dispersion for cosmetics, inks, dyes and other chemicals
• Destroy undesirable colloidal dispersion during water and sewage treatment
• Reduce cost of additives by calculating the minimum amount needed to achieve a
desirable effect, such as amount of flocculent added to water during water treatment
• Many other uses in mineral processing, ceramics manufacturing , pharmaceutical
production, etc..,

SURFACTANT
• Surfactants are compounds that lower the surface tension (or interfacial tension)
between two liquids, between a gas and a liquid, or between a liquid and a solid. Surfactants
may act as detergents, wetting agents, emulsifiers, foaming agents. Surfactant contains one
polar group which is called hydrophilic and non polar group which is called hydrophobic
Examples: Sodium stearate, sodium lauroyl sarcosinate and carboxylate-based fluorosurfactants
such as perfluorononanoate, perfluorooctanoate (PFOA or PFO).

SURFACTANT
Applications of surfactant:
Mechanism of the action of detergent surfactant
• Detergent, formulation of solution, formulation of emulsion , formulation of ointment and
formulation of shampoo

STRUCTURE OF ELECTRICAL DOUBLE LAYER
• There are three models formulated for the structure of electrical double layer
1. Helmholtz – Perrin model
2. Gouy – Chapman model
3. Stern’s model

HELMHOLTZ – PERRIN MODEL FOR EDL
According to this model
• The electrical double layer is a flat condenser i.e., one plate of which is connected
directly to the surface of a solid (a wall) whereas the other plate which carries
the opposite charge (counter ions) is in a liquid at a very small distance away from
the first plate

• Helmholtz considered the double layer consisting of two oppositely charged layers at a
fixed distance a part. The charge density on the two sheets are equal in magnitude but
opposite in sign exactly as in parallel plate condenser
• The potential in the Helmholtz layer is described by the Poisson’s equation

• The potential falls in a straight line shown in the diagram as the distance x
from the surface of the solid increases
Limitations
• This model for the electrical double layer does not explain electro kinetic
phenomena
• The thickness of the Helmholtz – Perrin double layer is very small and
approximates molecular dimensions but experimentally it was shown
that the thickness of the liquid layer that has adhered to the solid surface
is greater than the thickness of the Helmholtz double layer
x

According to this model
• The counter ions cannot concentrate only at the interface and form a mono ionic layer
but are scattered in the liquid phase at a certain distance away from the interface. i.e., by
the electric field of a solid phase which attracts an equivalent amount of oppositively charged
ions as close as possible to the wall and also by the thermal motion of ions as a result of
which counter ions scattered in the liquid phase
GOUY – CHAPMAN MODEL FOR EDL

• As the distance from the interface increases the intensity of the electric field gradually
decreases and counter ions of the double layer scatter more and more due to thermal motion,
as a result, the concentration of counter ions decreases and becomes equal to that of ions
which are in the bulk of liquids
• The potential drop follows not a straight line, but a curve shown in the diagram, because the
counter ions which compensate the surface charge are distributed non uniformly. The curve
drops steeper in places where there are more compensating counter ions and smoother in
place where there is a small amount of compensating counter ions

• The change in concentration of the counter ions near a charged surface follows the
Boltzmann distribution
Where
ni is the number of ions of type i per unit volume near the surface
ni
0 is the concentration far from the surface (bulk concentration)
Zi is the valency of the ion
e is the charge of the ion
ψ is the potential work required to bring an ion from infinity to a position

STERN’S MODEL FOR EDL
According to Stern
• The first layer of counter ions are attracted to the wall
under the action of both electrostatic and adsorption
forces. As a result a part of counter ions is retained by
the surface at a very close distance as fixed layer in
which the electric potential falls sharply like Helmholtz-
Perrin model
• The remaining counter ions are randomly distributed in
the diffused part of the double layer due to thermal
motion of ions in which the potential falls gradually like
Gouy-Chapman model
• The diagrammatic representation of electrical double
layer and the electric potential drop is represented in the
diagram shown below

STERN’S MODEL FOR EDL
• When electrolytes are introduced into the system , the diffuse layer will contract and an
increasing number of counter ions will appear into the adsorption layer and zeta potential
decreases and approaching zero
• When the system is diluted the diffuse layer expands and zeta potential increases.
• The distribution of ions in the electric double layer is strongly affected by the nature of
counter ions.
• If the counter ion have higher valency, the thickness of the diffused layer is reduced and
zeta potential become lower down

OVERVOLTAGE AND POLARIZATION
• For a given cell
Pt, H2 (1 atm ) / HCl (a = 1) // Pt/O2 (1 atm). The emf is 1.12 volts. If the opposing voltage
is slightly greater than 1.12 V . The cell reaction should get reversed and evolution of H2 should
start. But actually this is not true, and an opposing voltage of 1.7 volts is required.
“The excess of voltage over the voltage of the cell that is necessary to cause the reverse
reaction is termed as overvoltage and this phenomenon is known as Polarization ”
i.e., 1.7 – 1.12 = 0.58 = over voltage
i.e., ∆ϕ - ∆ϕ eq= 
where ∆ϕ is the actual potential difference (EMF of the cell)
∆ϕ eq is the theoretical potential difference
 is the overvoltage or overpotential

HYDROGEN OVERVOLTAGE
• The difference of potential at which hydrogen gas is actually evolved during the electrolysis
and theoretically at which it happens is known as hydrogen over-voltage
Hydrogen over-voltage at various metal
ELECTRODE HYDROGEN OVER VOLATEGE (VOLTS)
Platinized platinum 0.25
Gold 0.53
Silver 0.41

THEORIES OF OVERVOLTAGE
a) Combination of atoms to form molecules
According to this theory
• The formation of hydrogen molecules from hydrogen atoms is a rate determining step
H + H → H2(g)
• The over voltage may be due to aggregation of hydrogen atoms at cathode
• The variation of over voltage from one metal to another was due to their differing catalytic
effects on the rate of combination of hydrogen atom
i.e., the order of hydrogen over-voltage of various metal is
Pt < Pd < Fe < Ag
i.e., the order of catalytic effect is
Pt > Pd > Fe > Ag

THEORIES OF OVERVOLTAGE
b) Neutralization of charge (or) ion discharge
According to this theory
• The neutralization of hydrogen ion by an electron is the cause for over – voltage
• i.e., the electron cannot cross the barrier to discharge the ion unless it requires sufficient
energy to pass over the top

ELECTRON TRANSFER REACTIONS : THE MARCUS THEORY
Definition:
• Chemical reaction in which an electron transferred from the reduced form of a reactant
to the oxidized form of a reactant in solution is called electron transfer method
Examples:
1. MnO4
2- + MnO4
- → MnO4
- + MnO4
2-
2. [Fe(H2O)6]2+ + [Fe(H2O)6]3+ →
[Fe(H2O)6]3+ + [Fe(H2O)6]2+

MARCUS THEORY FOR ELECTRON TRANSFER REACTION
• The electron transfer should obey the Franck-Condon principle
i.e., the electron transfer takes place so rapidly that the atoms in the reactants and in the
solvent molecules do not have time to move in that instant
• The solvent dipolar molecules are partially oriented towards the ions and much more
oriented towards the more highly charged ions. The newly formed ions in electron
transfer reaction suddenly attain a new solvent environment
i.e., an appropriate redistribution of the orientation of the solvent molecules in the
vicinity of each ion needs to occur prior to the electron transfer

ELECTRON TRANSFER REACTION
• Consider electron transfer form a donor species D to an acceptor species A in solution
• First D and A must diffuse through the solution and collide to form a complete DA in
which the donor and the acceptor are separated by a distance r which is the distance
between the edges of each species
i.e., D + A ⇌ DA ; K1 =
[𝐷𝐴]
𝐷 [𝐴]
DA → D+A- ; rate = Ket [DA]
D+A- ⇌ D+ + A- ; Kv =
D+ [A−]
[D+A−]
Mechanism and kinetics of electron transfer reaction:

• The rate constant Ket is expressed as
Ket = k  e (-∆G# / RT)
where k is the transmission co-efficient
 is the vibrational frequency with which the activated complex approaches the
transition state
∆G# is the Gibbs energy of activation

• The main mechanism of electron between D and A for a given distance r is tunneling
through the potential energy barrier shown in the diagram
• Initially the electron to be transferred occupies the HOMO (highest occupied molecular
orbital) of D, and the overall energy of DA is lower that of D+A-
• As the nuclei rearrange to a configuration represented by q* in fig b the highest occupied
electronic level of DA and the lowest unoccupied electronic level of D+A- become
degenerate and electron transfer becomes energetically feasible (fig b)
• After an electron moves from the HOMO of D to the LUMO of A, the system relax to the
configuration represented by qo
p in fig c
Electron tunneling in electron transfer reaction:

• ∆G# is determined by the expression
∆G# = (∆G0
r+)2 / 4
Where
∆G0
r is the standard reaction Gibb’s energy for the electron transfer process
•  is the reorganization energy, i.e., the energy associated with molecular rearrangements
that must take place so that DA can assume the equilibrium geometry D+A-. These
molecular rearrangements include the relative orientation of the DA and the relative
orientation of the solvent molecules surrounding DA
• The potential energy surfaces of the reactant complex DA and the product complex D+A-
are shown in the diagram

• The electron transfer occur only after thermal fluctuations bring the geometry of DA to q*,
the value of the nuclear coordinate at which the two parabolas intersect
• The factor k is a measure of the probability that the system will convert from reactants
(DA) to product (D+A-) at q* by electron transfer within the thermally excited DA complex

KINETICS OF ELECTRODE REACTIONS
A reaction occurring at an electrode surface involves the following steps in succession
1. Diffusion of the reactants to the electrode
2. Adsorption of reactants on the electrode
3. Transfer of electrons to or from the adsorbed reactant species
4. Desorption of products from the electrode
5. Diffusion of products away from the surface of the electrode

• When an electrochemical cell operates under equilibrium the rate of electron transfer in the
anodic direction so that the current density (current per unit area) at cathode (ic)
i.e., is equal to current density at anode (ia)
ic = ia = i0
• The current density i0 at equilibrium is called the exchange current density. The rate r of a
chemical reaction at the surface of an electrode is given by
r = i/z F
• Where z is the charge on the ionic species
F is the Faraday 96458 C mol-1
i.e., For an electrochemical reaction, r  i

• When an electrochemical cell operates under non-equilibrium conditions, ic ≠ia and there is
net current density i = ia – ic In such a case the potential difference between the cell terminals
departs from the equilibrium value ∆ϕ = E, the cell EMF. If the cell is converting chemical
free energy into electrical energy, ∆ϕ < E. If on the other hand, the cell is using an external
source of energy to cause the chemical reaction, ∆ϕ > E. The actual value of ∆ϕ depends upon
current density i at the electrodes. It is customary to define the quantity over potential  of
the cell as
∆ϕ - ∆ϕ eq= 

• Kinetics of an electrode reaction (Butler volmer equation)
Consider the electrode reaction
Mz+
(aq) + ze → M(s)
• Which occurs when the reactants ion is in the vicinity of an
electrode surface so that the electrons are transferred from
the electrode to the ion
• According to eyring activated complex theory (ACT), the
rate constant k2 of the chemical reaction is given by
k2 = B exp (-∆G
≠
/ RT)
• Where ∆G
≠
is the Gibbs free energy of activation and B is
some constant
KINETICS OF ELECTRODE REACTIONS:BUTLER VOLMER EQUATION

• Consider a reaction at the electrode in which a particular species is reduced by the transfer
of a single electron in a rate-determining step. Let [Ox] and [Red] be the concentration of
the oxidized and reduced forms of the species, outside the double layer. Clearly the net
current at the electrode is the difference of the currents resulting from the reduction of Ox and
oxidation of red. The rates of these process are kc[ox] and ka[Red], respectively. In a reduction
process, the magnitude of charge transferred per mole of reaction events is
F = eNA
Where F is the faraday constant

• The cathodic current density ic, arising from the reduction is given by
ic = F kc [Ox]
• An opposing anodic current density ia, arising from oxidation is given by
ia = F ka [Red]
• Where kis are the corresponding rate constants. Hence the net current density at the electrode is
given by
i = ia – ic = F ka [Red] - F kc [Ox]
= F Ba [Red] exp (-∆Ga
≠
/ RT) – F Bc [Ox] exp (-∆Gc
≠
/ RT)
• When ia > ic so that i > 0, the current is anodic and ic > ia so that i < 0, the current is cathodic

• Let us consider a reduction reaction. As an electron is transferred from one electrode to
another, the electrical work done is e∆, where is the electrical charge and ∆ is the
potential difference between the electrodes.
• Hence the Gibbs free energy of activation is changed from ∆G
≠
to ∆G
≠
+ F ∆,
if the transition state corresponds to Ox being very close to the electrode.
• Thus, if ∆ > 0, more work has to be done to bring Ox to its transition state,
with the result that Gibbs free energy of activation is increased.
• On the other hand, if the transition state corresponds to Ox being far from the electrode
i.e., close to the outer plane of the double layer, then ∆G
≠
is independent of ∆

• In practice, however the situation is midway between the two extremes.
• Hence, we can write the Gibbs free energy of activation for reduction as ∆G
≠
+  F ∆,
where  , called the transfer coefficient or symmetry factor, lies between 0 and 1, i.e.,
0<  <1.
• Let us next consider the Oxidation of Red. Here Red discard an electron to the electrode
with the result that the extra work needed for reaching the transition state is zero if this state
lies close to the electrode.
• If that state lies away from the electrode (i.e., close to the outer plane of the double layer),
the work needed is - F∆ so that ∆G
≠
changes to ∆G
≠
- (1 -  )F ∆.

• Substituting the two Gibbs free energies of activation , we obtain the following expression
for the current density
• At equilibrium, ∆ = ∆eq and the net current is zero and the equilibrium current densities
are equal. Thus, if the potential difference differs from its equilibrium value by the
over potential, so that

• The two current densities are
• Since the two equilibrium current densities, ia,e and ia,c are equal, we can drop the
subscripts designate each of them as i0, the exchange current density and
• This is known as Butler Volmer equation

• The plot of current density (i) Vs. the over potential () in accordance with the Butler – Volmer
equation is shown in the diagram

• In the case A there are high exchange current density i0 at both electrodes (the individual
electrode curves are labeled as 𝐴′ and 𝐴′′ . In this case even a small over potential will
produce appreciable flow through the cell )
• The other case B corresponds to very low exchange current density i0 . In this case, a large
value of over potential is required to cause appreciable current flow through the cell

TAFEL EQUATIONS AND TAFEL PLOT
• We know that the Butler -Volmer equation is
--------------(1)
• When the over potential is large and positive, the second exponential in equation (1)
is much smaller than the first and may be neglected, giving
--------------(2)
• Taking ln on both side of equation (2) we get
ln i = ln i0 + (1-𝛼)F /RT --------------(3)

• When the over potential is large but negative the first exponential in equation (1) is
much smaller than the second and may be neglected and we have
--------------(5)
--------------(4)
• Taking ln on both side of equation (4) we get
ln i = - {ln i0 -𝛼F /RT}
ln i = - ln i0 + 𝛼F /RT
- ln i = ln i0 - 𝛼F /RT
ln (-i) = ln i0 - 𝛼F /RT

• Equation (3) & (5) are called Tafel equation
• The plot of the logarithm of the current density
(i) against the over potential shown in the
diagram is called Tafel plot
• The linear portion of the curve in this diagram
agree with the Tafel equation
• From the slop and the intercept of the Tafel
plot, 𝛼 and i0 can be determined

ELECTROCHEMICAL PASSIVITY
• We know that metals dissolve at anode producing cations
(oxidation)
i.e., M ⇌ Mn+ + ne- (anode)
• This will be happened only the potential applied is greater than
the reversible electrode potential of the metal. “On increasing the
current density , step by step, a stage is reached at which the
anode potential rises suddenly, but current strength drops, as
shown in the diagram and the metal almost ceases to dissolve .
This state is called Passive state and this phenomenon is called
passivity ’

CONCENTRATION POLARIZATION
• The phenomenon of the departure of the electrode potential (increase or decrease)
from the reversible value as a result of the change of concentration in the vicinity
of the electrode is known as concentration polarization and it is the basis of the
polarographic method of analysis (Dropping mercury electrode)

• For an anodic reaction (oxidation)
M → Mn+ + ne-
• Metal ionic concentration increases at the vicinity of the metal and the electrode potential
is increases as per the Nernst equation
E (M+, M) = E0
(M+, M) +
𝑅𝑇
𝑛𝐹
ln [Mn+ ]
• For a cathodic reaction (reduction)
Mn+ + ne- → M
• Metal ionic concentration decreases at the vicinity of the metal and the electrode potential
is decreases as per the Nernst equation
E (M+, M) = E0
(M+, M) +
𝑅𝑇
𝑛𝐹
ln [Mn+ ]
CONCENTRATION POLARIZATION

POURBAIX DIAGRAM [PE VS PH], [E0 VS PH]
[Electrode potential Vs PH]
• Graphical representation of the thermodynamics equilibrium state of a metal –electrolyte system
is called Pourbaix diagram
Eg: Iron Pourbaix diagram

• The lines of the diagram dividing different zones of the equilibrium states are calculated by
• Nernst Equation
E = E0 – (
0.059
𝑛
) ln Cion
• where, E0 is the standard electrode potential
0.0591/n is the number of electron transferred
Cion is the molar activity of ions

• Pourbaix diagram help us to determine the corrosion behavior of a metal in water solution.
i.e., the direction of the electrochemical process and the equilibrium state of a metal at a
certain electrode potential in a water solution at a certain value of PH
• Normally Pourbaix diagram built for the low concentration of metal ions
• Dashed line (blue colour) enclose the theoretical region of the stability of the water
2H2O → O2 + 4H+ + 4e-
• Above the dashed line, water is oxidized to oxygen
• Below the lower dashed line, water is reduced to hydrogen. 2H+ + 2e- →H2

Solid iron zone : (Below a-b-j)
(Immunity zone)
• The electrochemical reactions in this zone proceed in the direction of reduction of iron ions
• No corrosion occurs in this zone
Fe2+ zone : (Below a-b-n-c-d-e)
(corrosion zone)
• Aqueous solution of ion (Fe2+)
• Metallic ion oxidizes in this zone (Iron corrodes in this zone)

Fe3+ zone : (e-d-f-g-k)
(corrosion zone)
• Aqueous solution of ion (Fe3+ )
• Metallic ion oxidizes in this zone (Iron corrodes in this zone)
Fe2O3 (solid) zone : (e-d-f-h-i)
(Passivation zone)
• Solid ferrous oxide (Fe2O3)
• Iron corrodes in this zone however the resulted oxide film depresses the oxidation
process passivation
• Passivation means – Protection of the metal from corrosion due to the formation
of a film of a solid product of the oxidation reaction

Fe2O3 (solid) zone : (e-d-f-h-i)
(Passivation zone)
• Solid ferrous oxide (Fe2O3)
• Iron corrodes in this zone however the resulted oxide film depresses the oxidation
process passivation
• Passivation means – Protection of the metal from corrosion due to the formation
of a film of a solid product of the oxidation reaction
Fe3O4 [Fe2O3.FeO](Solid oxide) zone : (n-c-i-p)
(Passivation zone)
• The iron oxide film causes passivation

Fe(OH)2 (solid hydroxide) zone : (b-n-p-j)
(Green rust )
(Passivation zone)
• Green rust is an unstable corrosion product typically produced in low-oxygen environment
• Green rust occurs when the concentration of OH is more (PH increases )
Horizontal lines: (a-b), (e-d)
• Represent redox reactions which are independent of PH
Fe(s) → Fe(aq)
2+ + 2e-
Fe2+ → Fe3+ + e-

Vertical line : (d-f)(b-n)
• Represents non redox reaction, electrons are not involved which are dependent on PH
• (Acid base reaction)
2Fe3+
(aq) + 3O2-→ Fe2O3
Fe2+ + 2OH- → Fe(OH)2
Diagonal line : (c-d)(b-j)
• Represents the redox reaction, which are dependent on PH
• (Acid base reaction)
(c-d) 2Fe2+
(aq) + 3H2O→ Fe2O3(s) + 6H+
(aq) + 2e-
(b-j) Fe + 2OH- → Fe(OH)2 + 2e-

• Active metal Fe is stable below the H2 line this means that iron metal is unstable in
contact with water
• Water is stable only below the dashed line

EVAN’S DIAGRAM
• In a corroding system an oxidation and a reduction must take place.
• The Evan’s diagram shows the relationship of current and potential for
the oxidation and the reduction reaction. These are usually plotted as
potential versus the logarithm of the current (E-lg I) curves.

EVAN’S DIAGRAM
• If these two reactions are responsible for the corrosion and no currents flow from or
into the systems, all the electrons released by the oxidation must be accepted by the
reduction. So the two reactions can only corrode at the potential where both reactions
currents are the same. This means the curves of the two reactions should intersect in
the Evan’s Diagram. This means, if the intersection in the Evan’s Diagram of the two
reactions is know, the corrosion potential and corrosion current are known.
.

• The popular Tafel Analysis is based on these theories.
EVAN’S DIAGRAM

ELECTROCHEMICAL CORROSION
Definition:
• Destruction of a metal through an unwanted electrochemical attack is called
electrochemical corrosion
Types
1. Wet corrosion. Eg: Rusting of iron
2. Galvanic corrosion Eg: Corrosion of zinc contaminated with copper on exposing
an electrolyte

Mechanism for wet corrosion (Rusting of iron)
• The diagrammatic representation for the rusting of iron in neutral aqueous solution
in the presence of atmospheric oxygen is shown as

• The Fe 2+ ions and 2OH- ions diffused towards and precipitated as Fe(OH)2
i.e., Fe 2+ +2OH- → Fe(OH)2
• If enough oxygen is present Fe(OH)2 is easily oxidized to ferric hydroxide
4Fe(OH)2 + O2 + 2H2O → 4Fe(OH)3 [2Fe2O3.6H2O]
• The resulted ferric hydroxide is called rust and as formula Fe2O3.xH2O. the value of
x varied depending upon the availability of water
 At anode (Oxidation )
Fe → Fe 2+ + 2e-
 At cathode (Reduction )
½ O2 + H2O + 2e- → 2OH-

Mechanism for Galvanic corrosion:
• The diagrammatic representation for the corrosion of zinc which is contaminated with
copper on exposing an electrolytic solution (acidic solution)

i.e., Zinc (higher in electrochemical series) form the anode and is attacked and gets
dissolved; where as copper (lower in electrochemical series or more noble) acts as cathode
 At anode (Oxidation )
Zn → Zn 2+ + 2e-
 At cathode (Reduction )
2H+ + 2e- →H 2(g)
 Net charge
Zn + 2H+ → Zn 2+ + H 2(g)

• In this protection method, the metallic structure (to
be protected) is connected by a wire to a more
anodic metal, so that all the corrosion is
concentrated at this more active metal
• The more active metal itself gets corroded slowly
while the parent structure (cathodic) is protected
• The more active metal used in this method is called
sacrificial anode
Prevention of electrochemical corrosion
1.Sacrificial anodic protection method

• In this method, an impressed current is applied
in opposite direction to nullify the corrosion
current, and convert the corroding metal from
anode to cathode
• The impressed current is derived from a battery
with an insoluble anode like graphite
Prevention of electrochemical corrosion
2. Impressed current cathodic protection

• Power storage systems are rechargeable device which generate electrical energy
from solar arrays or fuels or electric grid
Examples :
1. Solar cells
2. Fuel cells
3. Batteries
POWER STORAGE SYSTEM

• Fuel cells are galvanic cells in which chemical energy of fuels is directly converted into
electrical energy
Example :
1.Hydrogen-Oxygen fuel cell
2. Hydrocarbon – Oxygen fuel cell
3. Coal – Fired fuel cell
FUEL CELLS

HYDROGEN OXYGEN FUEL CELL
• The schematic diagram of this cell is shown in the figure
• It consist of two electrodes made up of porous graphite impregnated with a catalyst (Platinum,
silver or metal oxide)
• The inner side of the graphite electrode are in contact with an aqueous solution of KOH or
NaOH
• Oxygen and Hydrogen are continuously supplied into the cell under a pressure of about 50 atm
• The gases diffused into the electrode pores and the electrolytic solution
• The half-cell reaction at the electrode are as follow

• Oxidation half – cell reaction :
Hydrogen is oxidized to H+ ions which are neutralized by the OH – ions of the electrolyte
H2 → 2H+ +2e-
2H+ + 2OH- → 2H2O
The net oxidation half cell reaction is
H2 + 2 OH- → 2H2O + 2e-

• Reduction half – cell reaction :
Reduction half cell reaction involves the reduction of oxygen to OH- ions
O2 + 2H2O + 4e- → 4 OH-
• The overall fuel cell reaction is
2H2 + O2 → 2H2O
• The emf of the cell is found to be one volt. The water produce vaporizes off since the cell is
operated at temperature above 100 0C. This can be condensed and used

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