Transcript Unit - II Electrochemistryx

CY6251-Unit – II
ELECTROCHEMISTRY AND CORROSION
Lecturer-II /I Electrochemical cells and its types.
Electrode potential –origin – oxidation and reduction potent
INTRODUCTION
• Electrochemistry is a branch of chemistry.
• It deals with the chemical reactions produced by passing electric current through an
electrolyte or production of electric current through a chemical reaction
Conductors
Conductor is a material which allows free flow of electricity.
Example: All metals, graphite, fused salts, solution of electrolytes
Non-conductors (Insulators)
Insulators are materials which donot conduct electrical current
Example: Wood, plastics, most of non metals.
Types of conductors
(i) Metallic conductors : The solid material, which conduct electric current due to the
movement of electron from one end to the other end without producing chemical
reaction.
Examples : All metals & graphite.
(ii) Electrolytic conductors : They conduct electric current due to the movement of
ions from one electrode to another electrode in solution or in fused state. This process is
accompanied by a chemical reaction.
Examples : Metal ions dissolved solvent
Cell Terminology
1. Current: Flow of electrons through a conductor.
2. Electrode: Electrode is a material (rod, bar, strip) which conducts electrons.
3. Anode: Electrode at which oxidation occurs.
4. Cathode: Electrode at which reduction occurs.
5. Electrolyte: Water soluble substance forming ions in solution and conducts electric
current
6. Anode compartment: Compartment of the cell in which the oxidation half reaction
occurs. It contains the anode
7. Cathode compartment: Compartment of the cell in which the reduction half
reaction occurs. It contains the cathode
8. Half–cell: It is the part of a cell, which contains an electrode dipped in an
electrolyte. If oxidation occurs in this half-cell, then it is called the oxidation half
cell. If reduction occurs at the cell, it is called the reduction half-cell.
9. Cell: Device consisting of two half cell. The two half cells are connected through
one wire.
10.Salt bridge: Contains solutions of a salt (KNO3 or NH4NO3) that literally serve as
a bridge to completed the circuit, maintain electro neutrality of electrolyte and
minimize. For precise measurement of potential a salt bridge is used.
TYPES OF CELLS
A cell is a device consisting of two half cells. Each half cell consists of an electrode
dipped in an electrolyte solution. The two half cells are connected through one wire.
S.No
Electrolytic cell
Electrochemical cell
1
Electrical energy converted to
Chemical energy converted to electrical
chemical energy.
energy.
Example: Electrolysis, electroplating. Example: Daniel Cell
2
3
4
5
6
Anode carries +ve charge.
Cathode carries – ve charge.
Electrons are supplied to the cell
from an external source.
Amount of electricity is measured by
coulometer
Extent of chemical change is
governed by Faraday’s laws.
Anode carries –ve charge
Cathode carries +ve charge.
Electrons are drawn from the cell
Emf produced is measured by
potentiometer
Emf of the cell depends on the
concentration of the electrolyte and the
nature of the electrode.
Electrolytic cell - Example : Electrolysis of HCl.
Mechanism
At anode : 2Cl–

Cl2 + 2e (oxidation)
At cathode : 2H+ + 2e–

H2 (reduction)
Electrochemical cell - Example : Daniel cell
Electromotive Force (emf)
(at anode)
(at cathode)
Cu2+ + Zn  Zn2+ + Cu
(net cell reaction)
• At anode : Oxidation of Zn to Zn2+ place with the liberation of electrons.
• At cathode : Reduction of Cu2+ to Cu place by the acceptance of electrons. The
electrons liberated in oxidation reaction flow through external wire and are
consumed by the copper ions at the cathode.
• Salt bridge : It consists of a U-tube containing a saturated solution of KCl or
(NH4)2NO3 agar–agar gel. It connects the two half cells.
Functions
i. It eliminates liquid junction potential.
ii. It provides a path for the flow of electrons between two half cells.
Representation (Cell diagram)
i. Galvanic cell consists of two electrodes, anode and cathode
ii. Anode is written on the LHS and cathode on RHS
iii. The anode is written with the metal first and then the electrolyte which are separated
by a vertical line Examples : Zn/Zn2+ (or) Zn/ZnSO4
iv. The cathode is written with the electrolyte first and then the metal.
Examples : Cu2+/Cu (or) CuSO4/Cu
v. The two half cells are separated by a salt bridge, which is indicated by two vertical
lines.
Cell is represented as Zn/ZnSO4 (1M) // ) CuSO4 (1M) /Cu
ELECTRODE POTENTIAL ORIGIN OF ELECTRODE POTENTIAL
When a metal (M) is placed in a solution of its own salt (Mn+) one of the two processes
are possible
(i) Metal atoms go into solution in the form of ions.
M  Mn+ + ne- (Oxidation)
Example :
e-
(ii) Metal ions from solution may deposit on the metal
Mn+ + ne-  M (reduction)
e-
Zn
Cu
Zn
Cu
2+ Zn2+ + Cu2+
Zn 
2++ 2e-
2eCu
Electrode Potential
At equilibrium, the potential difference becomes a constant value which is known as the
electrode potential of the metal. Thus the tendency of the electrode to lose electrons is called
Oxidation potential and tendency of an electrode to gain electrons is called reduction potential.
Single electrode potential (E) : It is the tendency of a metallic electrode to lose or gain electrons
when it is in contact with a solution of its own salt.
Standard electrode potential (Eo) : It is the tendency of a metallic electrode to lose or gain
electrons when it is in contact with a solution of its own salt of 1M concentration at 25oC.
Lecturer-II /II Measurement Of Single Electrode Potential and its applications.
Electro chemical series and its significances
MEASUREMENT OF SINGLE ELECTRODE POTENTIAL
It is impossible to evaluate the absolute value of a single electrode potential. Using
reference electrode.
Reference (or) Standard electrode
The potential of unknown electrode can be measured by coupling it with another
electrode, called reference electrode whose electrode potential is already known.
Examples : Standard hydrogen electrode, Standard calomel electrodes.
Standard hydrogen electrode (SHE)
It is also called as Primary reference electrode because. The potential developed by
this electrode is arbitrarily fixed as zero
Construction
• It consists of a platinum foil that is connected to a platinum wire sealed in a glass tube.
•The Pt foil is dipped in 1M HCl. H2 gas of 1 atm pressure is passed through the side of
glass tube.
• The standard electrode potential of SHE is arbitrarily fix as zero
Pt , H2 (1atm) / H+(1M) ; E0 = 0V
H 2 (g)  2H + 2e
Limitations (or) drawbacks of SHE
• It is difficult to get pure hydrogen gas.
• The pressure of hydrogen is to be kept 1 atm all
the time.
• It is difficult to set up and transport.
• Hydrogen gas reduces many ions like Ag+ and
affects compounds of Hg, Ag etc
• A large volume of test solution is required.
• It cannot be used in solutions of redox systems,
the solution may poison platinum surface.
Saturated calomel electrode (SCE) (Secondary reference electrode)
• Glass tube containing pure Hg at the bottom over which mercurous chloride is
placed. The remaining portion of the tube is filled with saturated solution of KCl.
• The bottom of the tube is sealed with a platinum wire. The side tube is used for
making electrical contact with a salt bridge.
Hg | Hg2 Cl2(s) | KC | (Saturated, Solution) Eº = 0.2422V
2Hg(l ) + 2Cl  Hg 2Cl2 (s )+ 2e
HgCl2 +2e  2 Hg +Cl
KCl
(v)
0.1N
0.3335 V
1N
0.281 V
Saturated
0.2422 V
Measurement of single electrode potential using a reference electrode (saturated
calomel electrode)
Hg 2 Cl 2 (s)  2Hg ( ) 2Cl
The emf of the cell is measured using a potentiometer. The value of Ecell = 1.0025 volt.
E
= E oright----- Eoleft
E
= E calo --- EoZn
E oZn
= E ocal ---- Eocell
= + 0.2422 – 1.0025 = – 0.7603 V
ELECTROCHEMICAL SERIES (OR) EMF SERIES
Significance of emf series (or) Application of electrochemical series
• Calculation of Standard emf of a Cell : We can calculate the standard emf of a cell,
if the standard electrode potential values are known (Ecell =ERHE –ELHE)
• Relative ease of oxidation or reduction
(a) Fluorine has higher +ve value of standard reduction potential (+2.87V) and
shows higher tendency for reduction.
(b) Lithium has highest – ve value (–3.02V) and shows higher tendency
towards oxidation.
• Displacement of one element by the other
Metals with a lower reduction potential will displace metals with a higher reduction
potential from their salt solution (Copper will displace silver from its solution).
Example :
Zn(-0.76V) will displace copper (+0.34V) from its solution
Zn + CuSO4ZnSO4 + Cu
• Determination of equilibrium constant (K) for a reaction
Standard electrode potentials are used to determine the equilibrium constants as
follows: Go = ln K 2.303 RT log K
log 𝐾 =
𝐺0
2.303 𝑅𝑇
=
𝑛𝐹𝐸0
[𝐺0𝑛𝐹𝐸0]
2.303 𝑅𝑇
• Hydrogen displacement behavior
Metals with negative reduction potential (metals placed above H2) in emf
series will displace hydrogen from dilute acids solutions.
Example: Zn (-0.76 V) will displace H2 from dilute acids whereas, silver (0.8)
cannot
• Zn + H2SO4  ZnSO4 + H2
• Ag + H2SO4  No reaction
• Predicting the spontaneity of redox reactions
If Eo of a cell is positive the reaction is spontaneous.
If Eo of a cell is negative the reaction is not feasible.
Lecturer-II / III NERNST EQUATION FOR ELECTRODE POTENTIAL
The potential of any electrode system (E) depends on
(i) nature of the metal / element
(ii) temperature and concentration of the electrolyte. The functional dependence of
potential of any electrode system (E) on these factors is given by Nernst equation
E = E 0 + (RT / nF) ( log Mn+)
Potential evolves from the inter-conversion between chemical and electrical forms of
energy. The equation pertaining to these two forms of energy is
G = G0 + RT ln K ----- (1)
where G, G0, K and R are respectively
The free energy change for any chemical to electrical energy conversion process is
given by the equation
G = - nFE ----- (2)
where E is the potential of the electrode system,
F – Faraday’s constant = 96496 or 96500 coulombs and ‘n’ is the number of electrons
transferred between the element and the ion which are in equilibrium.
Combining equations (1) and (2), we get
-nFE = - nFE0 + RT ln K ----- (3)
where E0 is the standard electrode potential of the electrode system.
Consider the reduction reaction Mn+ + n e → M
Equation (3) becomes E = E0 - (RT / nF) ln (aproducts / areactants ) ----- (4)
where ‘a’ is the activity of the species. The activity of a (homogenous / uniform) solid
is taken as unity, that of the electrolyte expressed in terms of the concentration and that
of a gas (or gaseous mixture) expressed in terms of pressure (or partial pressure) of the
gas.
Equation (4) can be written as
E = E0 - (RT / nF) ln (aM / aMn+ )
(or) E = E0 + (RT / nF) ln [Mn+] ----- (4)
where [Mn+] is the concentration of the electrolyte / metal ion in solution.
i.e. E = E0 + (RT / nF) ln [Mn+] ----- (5)
substituting the values of R= 8.314 Joules, F = 96500 Coulombs and introducing the
factor 2.303 to convert natural logarithms to common logarithms, equation (5) becomes
E = E0 + (0.059 / n) log [Mn+] ----- (6)
Equations (5) and (6) are the two forms of Nernst equation which gives the dependence
of electrode potential on the factors mentioned
Applications :
• It is used to calculate the electrode potential of unknown metal.
• To know the emf and polarity of electrodes in an electrochemical cell.
• The corrosion tendency of metals in a given set of environmental
• conditions can be predicted.
Lecturer-II /IV Corrosion-Introduction, types Mechanism of dry and wet corrosion
“Corrosion is defined as the gradual destruction of metals or alloys by the chemical or
electrochemical reaction with its environment.”
Causes of corrosion occurs
Most of the metals (except noble metals) naturally exist in combined form. During
metallurgy the metal are extracted from their ores by reduction process. In the pure
metallic state, the metals are unstable and considered to be in the excited (higher
energy) state. Therefore, the extracted metals have a tendency to go to
thermodynamically stable (lower energy) state, which is otherwise known as corrosion.
Thus, corrosion is a process “reverse of extraction of metals.”
Consequences of corrosion
• The efficiency of machine will be lost due to the loss of useful properties of metal.
• The products gets contaminated due to corrosion.
• Increase in maintenance and production cost.
• Preventive maintenance like metallic (or) organic coating is required.
• Toxic products are released.
TYPES OF CORROSION
Based on the environment, corrosion is classified into two types.
(i) Dry (or) Chemical corrosion
(ii) Wet or electro chemical corrosion
Dry (or) Chemical corrosion
•
It is due to the direct chemical attack on metals by atmospheric gases
such as oxygen, carbon-di-oxide, hydrogen sulphide, etc.
•
It follows adsorption mechanism. Corrosion product accumulate in the same spot of
corrosion.
Example : Tarnishing of silver in H2S gas, Action of dry HCl on iron surface.
There are three main types of chemical corrosion.
•
Oxidation corrosion
•
Corrosion by hydrogen
•
Liquid–Metal corrosion
Oxidation corrosion
2M 2Mn+ + 2ne–
At the oxide scale/environment interface
n O + 2ne–  nO2–
2
2
The over all reaction
2M + n2 O2  2Mn+ + nO2–
Metal – Oxide scale
The nature of oxide film formed on the metal surface plays an important role in
oxidation corrosion.
Nature of oxide layer and future course of corrosion
1. Stable : The oxide layer formed in some cases stick firmly to the parent metal
surface. Such layers naturally do not allow penetration of oxygen to the underlying
metal surface and thus act as protective films. Example : Al, Sn, Pb, Cu, etc.
2. Unstable : In case of some metals, elemental or uncombined state is naturally more
stable than the combined state such as oxide, sulphide, sulphate, etc.
Metal oxide  Metal + Oxygen
Only forward reaction is favoured. So, oxidation corrosion is not possible in those
metals. Example : Ag, Au and Pt
3. Volatile : In some metals the oxide layers formed are volatile. They leave the metal
surface as soon as they are formed. That means, the fresh metal surface is kept exposed
all the time for further attack. This makes the corrosion continuous and rapid.
Example : Mo (MoO3 is volatile).
4. Porous : The oxide layer formed in some cases are porous. Atmospheric oxygen
gets free access to underlying metal surface. Consequently corrosion goes non-stop till
the entire metal is converted into its oxide. Example : Alkali metals (Li, Na, Ka etc.)
Pilling–Bed worth rule
An oxide layer is protective (or) non-porous, if the volume of the metal oxide formed is
atleast as great as the volume of the metal from which it is formed.
Protective or non-porous : Example : Al, Sn, Pb, Cu, etc.
An oxide layer is non-protective (or) porous, if the volume of the metal oxide formed is
less than the volume of the metal from which it is formed.
VM
2
O n VM
Non-protective or Porous :
Example :Alkali metals : Li, Na, K, etc.
Alkaline earth metals : Mg, Ca, Si etc.
Corrosion by hydrogen
Atomic hydrogen (H) can more easily penetrate steel and other metals than molecular
hydrogen (H2), which is chemically more active.
(i) Hydrogen embrittlement : At ordinary temperatures, some reactions produce
atomic hydrogen which attack metals and reduce their strength.
Example :
Aqueous solution of H2S liberates atomic hydrogen at iron-surfaces.
Fe + H2S  FeS + 2H
(ii) Decarburization: At high temperatures, atomic hydrogen is produced by thermal
dissociation.
H2  H + H
4H + C (in steel)  CH4
The gas collects in gaps and voids and causes blisters and fissures. Consequently the
metal becomes weak.
By other gases : Gases like, SO2, CO2, H2S, F2 and Cl2 are also corrosive.
c) Liquid–Metal corrosion
This is brought about by chemical action of flowing liquid metals at high temperatures
over solid metal or alloy. The corrosion involves either dissolution of a solid metal by a
liquid metal or internal penetration of the liquid metal into the solid metal.
Example : Coolant (sodium metal) leads to corrosion of cadmium in nuclear reactor.
Wet (or) Electrochemical Corrosion
This type of corrosion occurs when, Metal is in contact with electrolyte or varying
concentration of oxygen.
Two dissimilar metals or alloys are immersed or partially dipped in a solution.
At anode – Oxidation takes place (loss of electron)
At Cathode – reduction takes place (gain of electron)
i.e. It involves flow of current between anodic and cathodic areas.
Mechanism of electrochemical corrosion
Acidic : 2H+ + 2e–  H2
Basic: 2 O2 + 2e– + H2O  2OH–
The Overall reaction is
Fe + 2H+  Fe + H2
ii) Absorption of oxygen (Rusting of Iron) type corrosion
This type of corrosion occurs when metal comes in contact with neutral or alkaline
medium.
Over all : Fe2+ + 2OH–  Fe(OH)2 (Ferrous hydroxide)
4Fe(OH)2 + O2 + 2H2O  4 Fe(OH)3 (Ferric hydroxide)(Rust)
TYPES OF ELECTROCHEMICAL CORROSION
1. Galvanic (or) Bimetallic corrosion
In Zn–Cu couple, Zn acts as anode [E oZn 0.76V] and Copper acts as cathode [E ored
0.34V] (e.g.) Steel screws in a brass marine hardware.
2. Differential aeration (or) Concentration cell corrosion
This type of corrosion occurs when metal is exposed to varying concentration of
oxygen or an electrolyte.
At anode (less aerated part) corrosion occurs
M  M2+ + 2e–
At cathode (more aerated part) OH– ions are produced.
1 O + H O + 2e–  2OH–
2 2
2
Pitting corrosion
Pitting is a localised attack, resulting in the formation of a hole around which the metal
is relatively unaffected. The mechanism involved is differential aeration corrosion.
Fe  Fe 2 2e [Anode]
1
2
O2 + 2e + H2O  2OH [Cathode]
Fe +1/2 O2 + H2O  Fe(OH)2 [Over all Reaction]
4) Water line corrosion
Difference between Chemical and Electrochemical corrosion
S.No
1.
2.
3.
4.
5.
Chemical or Dry corrosion
Electrochemical or Wet corrosion
It occurs in the presence ofmoisture or
It occurs in dry condition
electrolyte.
It involves the direct chemical
It involves the setting up of large number of
attack of metalsby dry gases.
electrochemical cells.
Corrosion products accumulatein Corrosion occurs at the anode, products
the same spot of corrosion.
gather at the cathode.
Corrosion is uniform throughout Corrosion is not uniform, which depends on
the surface
anodic and cathodic.
It follows mechanism of electrochemical
It follows adsorption mechanism. reactions
Lecturer-II / V FACTORS INFLUENCING THE RATE OF CORROSION
The rate and extent of corrosion depends on
1.
2.
Nature of the metal
Nature of the environment
1. Nature of the metal
a.
Position in galvanic series
b.
Over voltage
c.
Purity of metal
d.
Relative areas of anode and cathode
e.
Nature of oxide film (surface film)
f.
Nature of corrosion product
g.
Physical state of metal
a) Position in galvonic series
(b) Over voltage
If hydrogen overvoltage of the cathodic metal is very high, corrosion rate is lesser.
When the voltage in a circuit of part of it raised above its upper design limit - over
voltage. 1/ corrosion (i.e.) Over voltage of a metal
(e.g.) When Zn dipped in 1 N H2SO4, the hydrogen over voltage is 0.70 volts.
c) Purity of metal: The 100% pure metal will not undergo any types of corrosion.
e.g. %Purity of Zn
Corrosion rate
99.9999
1
99.99
99.95
2650
5000
(d) Relative areas of anode and cathode : Corrosion is severe, if the anodic area is
smaller than cathodic area.
(e) Nature of oxide film (surface film): The nature of surface film formed decides the
extent of corrosion.
(f) Nature of corrosion product
The corrosion product formed is soluble in corroding medium, the corrosion rate will be
greater. Similarly, if the corrosion product is volatile
(e.g.) MoO3, SnCl4 both are volatile corrosion rate will be faster.
(g) Physical state of metal
The rate of corrosion depends on physical state of metal (i.e) grain size, stress,
orientation of crystals.
The smaller the grain size of metal or alloy greater will be its solubility and corrosion.
Areas under stress (bends, joints, rivets) tend to be anodic and leads to corrosion.
Nature of the environment
(a) Temperature: The rate of corrosion is directly proportional to temperature.
(b) Humidity: The greater is the humidity, the greater is the rate and extent of
corrosion. This is due to the fact that moisture acts as a solvent for O2, H2S, SO2 and
NaCl etc., resulting an electrochemical cell.
(c) Presence of impurity in atmosphere: Presence of gases like CO2, H2S, SO2 and
fumes of HCl, H2SO4 makes the environment acidic so corrosion is accelerated.
(d) Presence of suspended particles: If the environment contains chemically active
particles (NaCl, (NH4)2SO4) they absorb moisture and acts as strong electrolytes. So
rate of corrosion is high.
(e) Nature of ions present: Presence of silicate (anion) in the medium leads to
insoluble products (silica gel), which prevents corrosion.
Presence of Cl ions in electrolyte destroys the protective films and enhances corrosion.
If ammonium salts are present in corroding medium, that will lead to corrosion of many
metals.
(f) Effect of pH: Generally acidic medium is more corrosive than alkaline medium.
The corrosion of iron in oxygen free water is very slow upto pH = 5. But in presence of
oxygen the corrosion rate of iron is very high at pH = 5. But at pH = 4 the corrosion rate
of iron is considerably increased due to the change in oxidation states of iron from Fe2+
to Fe3+. Zinc which is readily corroded in acidic solution suffers very less corrosion in
alkaline medium i.e. pH = 11. Aluminium has less corrosion at pH = 5.5 which corrodes
rapidly at pH = 8.5.
Lecturer-II /VI Corrosion control-selection and designing of materials, Sacrificial
and Imparessed current cathodic protection.
The corrosion process depends on the metal and the environment. So control measures
are aimed at the metal and environment.
(i.e.)
(i) Metal based controlled measures and
(ii) Environment based controlled measures.
Metal based controlled measures
1.
Proper choice (or) Selection of metal
2.
Proper design
3.
Cathodic protection
Proper choice (or) Selection of metal
• Before fabrication of equipments, proper selection of metal or alloy have to be
studied according to the environmental conditions.
• Corrosion can be avoided by replacing the metal parts by plastics, rubbers, ceramics
etc. as far as possible.
Metals
Not affected by
Stainless steel
HNO3
Steel
Con.H2SO4
Ti
Hot acids and alkalies
Proper design
Design play an important role in corrosion control. Important design principle are
• Avoid direct contact of dissimilar metals by using adhesives.
• When two dissimilar metals are in contact, the anodic part should have large area than
the cathodic part.
• Care must be taken in designing an equipment so that the liquid can be drained off
completely.
• Crevice corrosion and caustic embrittlement can be prevented by avoiding hair
cracks, gaps and protruding parts (rivets etc).
• Single crystalline solids avoid inter granular corrosion.
• Stress corrosion can be avoided by heat treatment technique (annealing).
• Heterogeneity is reduced by using pure metals, which decreases corrosion rate.
(c) Cathodic protection
In cathodic protection, the corroding metal is forced to act as cathode.
• Sacrificial anodic protection
• Impressed current cathodic protection
Advantages
(i) Sacrificial anodic protection
No need of external power
supply.
Installation and maintenance
cost is low.
Limitations
Not suited for large objects
due to limited driving
potential.
Uncoated parts cannot be
protected.
(ii) Impressed current cathodic protection
In this method an impressed current is applied in the opposite direction of the corrosion
current to nullify it, which converts corroding metal from anode to cathode.
Advantages
• Larger objects can be
protected due to larger
driving voltage.
• Uncoated parts can be
protected.
Limitations
• Larger installation cost.
• Higher maintenance cost.
Lecturer-II /VII Paints-Constitution and functions
PROTECTIVE COATINGS: Protective coatings are the physical barrier between the
metal surface and the environment which prevents corrosion. They are also used for
decorative purpose and to impart some special properties such as hardness, electrical
properties, oxidation resistance, and thermal insulation.
Paints: It is a mechanical dispersion of one or more finely divided pigments in a
vehicle of medium. The medium consists of non volatile, film forming materials like
drying oils or resins in suitable volatile solvents (thinners).When paint is applied to a
metal surface, the thinner evaporates and the vehicle undergoes slow oxidation forming
a pigment film
Requisites (characteristics) of a good paint
• It should spread easily on the metal surface
• It should adhere well to the surface
• It should not crack on drying
• It should give stable and decent colour
• It should be corrosion and water resistant
Constituents of a paint and their functions
a)Pigments :
Pigments are solid and colour producing substances in the paint
Sl.No.
Colour
1.
2.
White pigments
Black pigments
3.
4.
5.
Red pigments
Blue pigments
Green pigments
Compounds used
White lead (2PbCO3.Pb(OH)2
Lithophone (75% BaSO4 ; 25% ZnS), TiO2
Lamp black, carbon black
Venetian red (Fe2O3 and CaSO4)
Indian red (Fe2O3)
Prussian blue Fe4[Fe(CN)6]
Chromium oxide
Functions
• It provides colour and opacity to the film
• It gives strength to the film
B ) Vehicle or drying oil
• It is a non volatile portion of the medium. It is a film forming constituent of the paint.
Examples : Linseed oil, dehydrated castor oil.
Functions
• They form protective film by the oxidation and polymerization of the vegetable oil.
• They hold pigment particles together on the metal surface.
• They impart water repellency, toughness and durability to the film.
c) Thinners ( solvents )
It is a volatile portion of the medium which easily evaporates after the application of the
paints.
Examples : Acetone, Diethylether, Chloroform, etc.
Functions
It reduces the viscosity of the paint
It acts as a dispersing medium for the oil and pigments
It increases the elasticity of the film
It increases penetrating power of the vehicle
d) Extenders (Fillers): They are low refractive indices materials which increases the
durability of paint.
Examples : Talc, gypsum, china clay.
Functions
• It reduces the cost of the paint.
• It retards the settling of the pigment and cracking.
e)Driers: They are oxygen carrier catalyst, which accelerate the process of drying.
Examples : Metallic soaps, linoleates, resinates of Co, Mn and Pb. Function
• They accelerate the drying of the oil film through oxidation, polymerisation and
condensation.
f) Plasticisers : These are chemicals added to paint to increase the plasticity and prevent
cracking of the film.
Examples : Triphenyl phosphate, tricresylphosphate, etc.
g) Anti skinning agents: These are chemicals that prevents gelling and skinning of the
paint film.
Example : Polyhydroxy phenol
Pigment volume concentration ( P.V.C )
It is a property of a paint which indicates the durability, adhesion and consistency.
Higher the value, lower will be these qualitie
PVC =
Volume of pigment in the Paint
Volume of (pigment + vehicle) in the paint
Lecturer-II /VIII Electoplating – introduction and method of electroplating
copper
Electroplating (Electrodeposition ): The process of deposition of coating metal on the
base metal by passing a direct current through an electrolytic solution which contains
the soluble salt of the coating metal.
Electroplating on metals
• To increase the resistance to corrosion of the
plated metal.
• To improve the hardness and physical
appearance.
• To increase the decorative and commercial
values of the metal.
• To increase resistance to chemical attack.
• To improve the surface properties.
On non-metals
• To increase strength.
• To decorate surfaces of non
metals like wood, plastic, glass.
• For obtaining surface
conductivity.
Electrodeposition: (Also knows as electroplating)
Electroplating is using a small sheet of metal in an electrocytic cell to coat another
object. It is used to protect objects from damage against rusting and corrosion of
metals.
Lecturer-II /IX Electroless plating-introduction and method of plating nickel.
Principle
Electroless plating is a technique of depositing a noble metal (noble metal salt solution)
on a catalytically active surface of the base metal by using a suitable reducing agent
without electric current. Metal ions + reducing agent  Metal + Oxidised products
1. Electroless nickel plating
Step 1 : Pretreatment and activation of the surface: The surface to be plated is
degreased first by using an organic solvent or alkali followed by acid treatment.
Step 2 : Plating bath
Sl.No.
Nature of the content
Compound
Quantity (g/L)
1.
Coating solution
NiCl2
20
2.
Sodium hypophosphite
20
3.
Reducing agent
Complexing agent cum
exhaltaant
4.
Sodium succinate
15
Buffer
Sodium acetate
10
5.
pH
4.5
-
6.
Temperature
93oC
-
Step 3 : Plating Procedure
The pretreated object is placed in the plating bath for the required time. The reduction
reaction occurs and nickel gets coated over the object
A) H2PO2- + H2O  H2PO3- + 2 H+ + 2eB) Ni2++ 2e-  Ni
Over all reaction Ni2+ + H2PO2- + H2O  Ni + H2PO3- + 2 H+
Applications
• It is used in domestic and automotive
fields (eg., jewellary, tops of perfume
bottle)
Advantages of electro less plating
over electroplating
• No electricity is required
• It can be easily plated on insulators
• Electroless nickel coated polymers are
used as decorative material.
• Complicated parts can be plated
uniformly.
• Copper and nickel coated plastic
cabinets are used in digital and electronic
instruments.
• This
coating
possess
unique
mechanical, chemical and magnetic
features