Chapter: Chemistry - Electro Chemistry and Corrosion

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Corrosion

1 Corrosion 1.1 Consequences of corrosion 2 Types of Theories of Corrosion 2.1 Dry or Chemical corrosion 2.2 Wet or electrochemical corrosion 3 Factors influencing the rate of corrosion 3.1 Nature of the metal 3.2 Nature of the environment 4 Corrosion Control 4.1 By modifying metal 4.2 BY modifying the environment 5 Paints 5.1 Characteristics of a good paint 5.2 Constituents and their functions of paints 6 Metallic Coatings 6.1 Electroplating 6.2 Electroless Plating


CORROSION


1 Corrosion

1.1 Consequences of corrosion

2 Types of Theories of Corrosion

2.1 Dry or Chemical corrosion

2.2 Wet or electrochemical corrosion

3 Factors influencing the rate of corrosion

3.1 Nature of the metal

3.2 Nature of the environment

4 Corrosion Control

4.1 By modifying metal

4.2 BY modifying the environment

5 Paints

5.1 Characteristics of a good paint

5.2 Constituents and their functions of paints

6 Metallic Coatings

6.1 Electroplating

6.2 Electroless Plating

 

 

1 Corrosion

 

It is the gradual deterioration of metals by chemical, electrochemical or biochemical interaction with the environment.

 

Causes of Corrosion

 

Metals occur in nature as their oxides, sulphides carbonates etc. The chemically combined state is thermodynamically more stable. When we extract a metal from its ore, the metal is in a higher energy state, which is thermodynamically unstable. So it tries to go back to the stable state by chemical or electrochemical interaction with the environment.

 

 

            Consequences or effects of Corrosion

 

            Efficiency of the machine decreases.

            Plant has to be shut down due to failure.

            Product is contaminated.

            The toxic products of corrosion cause health hazards.

            There is a necessity to over design to allow for corrosion.

 

 

2 Types  or Theories of Corrosion

 

I.  Dry or Chemical Corrosion

II. Wet or Electrochemical Corrosion

 

2.I. Dry or Chemical Corrosion

 

It is due to the attack on metal surface by atmospheric gases like O2, SO2, H2S etc. (e.g.) Tarnishing of silver by H2S.

 

There are three types of dry or chemical corrosion.

 

Oxidation Corrosion

Corrosion by Hydrogen

Liquid Metal Corrosion

 

Oxidation Corrosion

 

It is due to the direct attack of oxygen on metal surface in the absence of moisture. Alkali and Alkaline earth metals are corroded at low temperatures. At high temperatures, most metals except Au, Pt and Ag are oxidized.

 

Mechanism

                      Oxidation occurs at the surface of the metal to form M2+ ions.

          ® M2+ + 2e-

                      Oxygen takes up the electrons. O2 is reduced to O2-

        O2 + 2e-  ® O2-

                      O2- ion reacts with M2+ to form metal oxide.

M2+ + O2- ® MO

 

The metal surface    is converted to a monolayer of metal oxide. Further corrosion occurs by diffusion of M2+ ion through the metal oxide barrier. The growth of oxide film is perpendicular to the metal surface.



Different types of oxide films are formed.

(i) Porous and Non-Porous Oxide Film (or) Pilling-Bedworth Rule

 

(a) If the volume of the oxide layer formed is less than the volume of the consumed, the oxide layer is porous. (e.g.) The volumes of oxides of alkali and alkaline earth metals are less than the volume of the metal consumed. So the oxide layer is porous and non-protective

 

(b) If the volume of the oxide layer formed is greater than the volume of the metal consumed, the oxide layer is non-porous.(e.g.) The volumes of oxides of heavy metals such as Pb, Sn are greater than the volumes of the metal consumed. So the oxide layer is non-porous and protective.

 

(ii) Stable Oxide Layer

 

A stable oxide layer is firmly adsorbed on the metal surface. The layer is impervious and prevents further corrosion. So the layer itself acts as a protective coating. (E.g.) Oxides of Al, Cu etc.

 

 

(iii) Unstable oxide Layer

 

This is mainly produced on the surface of noble metals such Ag, Au etc. The unstable oxide decomposes to stable metal and oxygen. Metal Oxide Metal + Oxygen

 

(iv) Volatile Oxide

 

The oxide film volatilizes as soon as it is formed. It leaves fresh metal surface for further continuous attack. (e.g.) Molybdenum oxide MoO3.

 

(2) Corrosion by Hydrogen

 

Hydrogen embrittlement Definition

 

It is formation of cracks and blisters on the metal by hydrogen gas when the metal comes into contact with H2S. Iron liberates atomic hydrogen by reacting with H2S.

 

Fe + H2S   FeS + 2H

 

Hydrogen atoms diffuse into the voids of metal matrix. When the pressure of the gas increases, cracks and blisters develop on the metal.

 

            Decarburisation

 

It is the process of decrease in the carbon content of steel. At high temperature, molecular hydrogen decomposes to atomic hydrogen. High Temperature


When steel is exposed to this environment, carbon in the steel reacts with atomic hydrogen. 

C + 4H - - > CH4

Hence the carbon content in steel decreases. Collection of methane gas in the voids of steel develops high pressure and causes cracking.

 

(3) Liquid Metal Corrosion

It is due to the chemical action of flowing liquid metal at high temperature. It involves :

 

          dissolution of a solid metal by the liquid metal.

            Penetration of liquid metal into the solid metal.

 

2. 2. Wet (or) Electrochemical Corrosion :

 

It occurs under the following conditions.

 

When two dissimilar metals or alloys are in contact with each other in presence of an aqueous solution or moisture.

 

            When the metal is exposed to an electrolyte with varying amounts of oxygen.

 

Mechanism of Wet Corrosion

            Metal dissolution occurs at the anode.

M → Mn++ + ne-

            Reduction reaction occurs at the cathode in different environments.

 

Acidic environment : Here hydrogen gas is evolved at the cathode.

2 H+ + 2e- → H2

(b) Neutral environment : In neutral or slightly alkaline medium, hydroxide ions are formed at the cathode.

 

½ O2 + 2e- + H2O → 2OH-

 

(a) Hydrogen Evolution type corrosion (In Acidic Medium)

All metals above hydrogen in the electrochemical series tend to get dissolved

 

in acidic solution with simultaneous evolution of H2 gas. e.g.) When iron comes into contact with non-oxidising acid like HCl, hydrogen evolution occurs.

 

At anode : Iron is oxidized to Fe2+

 

Fe → Fe+2 + 2e-

 

At cathode : H+ ion is reduced to H2.

 

2 H+ + 2e- → H2

 

oxygen, OH- ions are formed.

 

At anode : Iron is oxidized to Fe


 

(b) Absorption of Oxygen (or) Formation of hydroxide ion type corrosion (In neutral or weakly alkaline medium)

 

The surface of iron is normally coated with a thin film of iron oxide. But if some cracks develop on the film, anodic areas are created on the surface. The rest of the metal part acts as cathode.(e.g.) When iron is in contact with an electrolyte solution in presence ofoxygen, OH- ions are formed. At anode : Iron is oxidized to Fe+2

 

At cathode : Production of OH- ions (more aeration) ½ O2 + 2e- + H2O → 2OH

 

Waterline corrosion

 

Let us consider metal tank partially filled up with water. The metal area above water line is exposed to higher concentration of oxygen (cathode) than the metal below water level. The metal less exposed to O2 acts as anode and corrodes. This is called water line corrosion.

 

Examples of differential aeration corrosion

       Pitting or localized corrosion

          Crevice corrosion

            Pipeline corrosion

 

iv) Corrosion on wire fence

 

(i) Pitting Corrosion

 

It is the localized attack resulting in the formation of a hole due to corrosion. Example : Metal area covered by a drop of water, sand, dirt etc.

 

The area covered by the drop or dirt acts as anode and corrodes. Theuncovered area exposed to air or O2 acts as cathode.The rate of corrosion is more if the cathodic area is larger and anodic area is smaller. Thus more material is removed from the same area and a pit is formed.

 

At anode : Iron is oxidized to Fe+2

Fe → Fe2+ + 2e-

 

At cathode : O2 is reduced to OH-. ½ O2 + H2O + 2e- → 2OH Overallreaction :

 

Fe2+ + OH- → Fe(OH)2

 


At anode : Iron is oxidized to Fe+2

Fe → Fe2+ + 2e-

 

At cathode : O2 is reduced to OH-. ½ O2 + 2e- + H2O → 2OH

 

Overall Reaction

Fe+2 + 2OH- → Fe(OH)2

 

If enough oxygen is present, Fe(OH)2 is oxidized to Fe(OH)2.

4Fe(OH)2 + O2 + H2O → 4Fe(OH)3

 

Differences between chemical corrosion and electro-chemical corrosion: Chemical Corrosion Electro-chemical Corrosion

        It occurs in dry condition It occurs in presence of moisture or electrolyte.

 

It occurs by the direct chemical attack on the metal by the environment.It occurs by the formation of a large number of anodic and cathodic areas.

 

Even a homogenous metal surface is corroded. Only heterogeneous or bimetallicsurface is corroded.

 

Corrosion products gather at the place of corrosion. Corrosion occurs at the anode, while the products form elsewhere.

 

        It is a self controlled process It is a continuous process

 

It takes place by adsorption mechanism. It follows electrochemical reaction.(e.g.) Mild scale formation on iron surface (e.g.) Rusting of iron under moist atmosphere

 

Types of electrochemical corrosion

There are two types:

(i) Galvanic corrosion

(ii) Differential aeration or Concentration cell corrosion

 

(i) Galvanic corrosion

 

When two different metals are in contact with each other in presence of aqueous solution or moisture, galvanic corrosion takes place.


The metal with more negative electrode potential acts as anode. Metal with less negative electrode potential acts as cathode. In the Zn-Fe couple as shown in the figure, zinc with more negative electrode potential, dissolves in preference to iron. Zn acts as anode and Fe as cathode.

 

Example :

 

Steel screw in a brass marine hardware easily undergoes corrosion. Iron has E0 = -0.44V. For Cu E0 = +0.34 V. Iron corrodes in preference to Cu.

 

Prevention

Galvanic corrosion is minimized by providing an insulation between the two metals.

 

(ii) Differential aeration (or) concentration cell corrosion


Let a metal be partially immersed in a conducting solution. The part of the metal above the solution is more aerated and acts like cathode. The less aerated metal part inside the solution acts as anode and corrodes.

 

At anode : Corrosion occurs (less aeration)

 

M → M2+ + 2e

 

At cathode : Production of OH- ions (more aeration) ½ O2 + 2e- + H2O → 2OH

 

Wateline corrosion

 

Let us consider metal tank partially filled up with water. The metal area above water line is exposed to higher concentration of oxygen (cathode) than the metal below water level. The metal less exposed to O2 acts as anode and corrodes. This is called water line corrosion.

 

Examples of differential aeration corrosion

       Pitting or localized corrosion

          Crevice corrosion

            Pipeline corrosion

           Corrosion on wire fence

 

(i) Pitting Corrosion

 

It is the localized attack resulting in the formation of a hole due to corrosion. Example : Metal area covered by a drop of water, sand, dirt etc.

 

 

The area covered by the drop or dirt acts as anode and corrodes. The uncovered area exposed to air or O2 acts as cathode. The rate of corrosion is more if the cathodic area is larger and anodic area is smaller. Thus more material is removed from the same area and a pit is formed.

 

At anode : Iron is oxidized to Fe+2

 

Fe → Fe2+ + 2e-

 

At cathode : O2 is reduced to OH-. ½ O2 + H2O + 2e- → 2OH Overall reaction :

 

Fe2+ + OH-→ Fe(OH)2

 

(ii) Crevice Corrosion

Let a crevice or crack between two different metallic objects be in contact

 

with a liquid. The crevice acts like anode due to less oxygen availability and corrodes. The exposed area acts as cathode.

 

(e.g.) rivets, joints.


 

(iii) Pipeline Corrosion

 

Buried pipelines or cables passing from one type of soil (clay, less aerated) to another type (sand, more aerated) get corroded due to differential aeration.

 

(iv) Corrosion on wire fence

 

In a wire fence, the wires at the crossings are less aerated than the rest of the

 

fence. So corrosion takes place at the wire crossings, which become anodic.

 

 

 

3     Factors influencing corrosion

 

3.1 Nature of the metal

(i) Position in emf series

 

Metals above hydrogen in the electrochemical series corrode easily because they have negative reduction potential. When two metals are in contact, the more active metal with a higher negative potential corrodes.

 

(ii) Areas of anode and cathode

 

Corrosion will be severe if the anodic area is smaller and cathodic area is larger. The larger cathodic area demands more electrons. So the anodic area corrodes faster.

 

(iii) Purity

 

100% pure metal will not corrode. (e.g.) Pure Zn does not corrode. If the metal has trace amount of impurity, it corrodes. (e.g.) Zinc metal with iron or copper impurity forms an electrochemical cell. The base metal Zn acts as anode and corrodes.

 

(iv) Over Voltage

 

Corrosion rate is inversely proportional to the over voltage of the metal in a corrosive surroundings. (e.g.) The hydrogen over voltage of Zn in 1M H2SO4 is 0.7V. So the rate of corrosion is low. But when some Cu impurity is present, the over voltage is reduced and corrosion rate increases.

 

(v) Nature of the Film

 

Nature of film formed on the metal surface determines extent of corrosion.(e.g.) In the case of alkali and alkaline earth metals, the oxide film formed is porous .The corrosion continues. In the case of heavy metals, the oxide film is non-porous. The film acts as a protective layer.

 

(vi) Nature of corrosion product

 

If the corrosion product is soluble in the corroding medium, corrosion rate is faster. Similarly if the corrosion product is volatile (e.g. MoO3), corrosion will be more.

 

3.2 Nature of Environment

(i) Temperature

Increase of temperature increases corrosion rate because the rate of diffusion of ions increases.

 

(ii) Humidity

 

Rate of corrosion is more, if humidity of environment is high. Moisture acts as solvent for O2, CO2 etc, to produce electrolyte necessary for formation of corrosion cell.

 

(iii) Corrosive gases

Acidic gases like CO2, SO2, H2S etc, produce electrolytes and increase corrosion.

 

(iv) Presence of suspended particles

 

Particles like NaCl, (NH4)2SO4 along with moisture are powerful electrolytes and increase rate of corrosion.

 

(v) Effect of pH

Generally in alkaline medium, the rate of corrosion is less compared to acidic medium.

 

The effect of pH on the corrosion of iron in water is shown in the Pourbaix diagram as indicated in the figure.

 


 

The figure shows zones of corrosion, immunity and passivity. Z is the point at which pH=7 and corresponding electrode potential is E= -0.2V. This is in the corrosion zone. So iron rusts under these conditions.

 

The rate of corrosion can be altered by shifting the point Z to different regions.

 

If the potential is changed to -0.8V by applying external current, iron becomes immune to corrosion.

 

         If the potential applied is positive, iron becomes passive.

         If the pH is increased to more than 7, corrosion rate decreases.

         If the pH is reduced to less than 7, rate of corrosion increases

 

 

4 CORROSION CONTROL

 

The rate of corrosion can be controlled by modifying the metal or environment. Some control methods are

 

         proper selection of metals

         Use of pure metals

         Use of metal alloys

         Cathodic protection

        Sacrificial anode protection

 

Impressed current cathodic protection 5) Changing the environment

 

6) Use of inhibitors

        Anodic inhibitors

        Cathodic inhibitors

7) Applying protective coatings

 

1) Proper selection of metals

 

Noble metals are used in ornaments and in surgical instruments, because they do not corrode. Contact of dissimilar metals far away from each other in electrochemical series should be avoided.

2) By using pure metals

 

Pure metals have high corrosion resistance. Even a trace of impurity will lead to corrosion, the base metal becoming anode.

 

3) Use of alloys

 

Use of metal alloys is a good method of protection against corrosion. (e.g.) Stainless steel containing chromium forms a coherent oxide film which protects steel against further attack

 

4.1 Proper designing

      Complicated designs with more angles, sharp edges and corners should be avoided.

 

Direct contact of dissimilar metals lead to galvanic corrosion. So insulating material between the two metals should be inserted.

 

           Smaller area for cathode and larger area for anode must be provided.

          Tanks and containers should be designed such that the liquid should be drained off completely.

        Crevices should be avoided or they should be filled using fillers.

          Bendings should be smooth.

Annealing minimizes corrosion.



5) Cathodic Protection

The metal to be protected is made to act like a cathode. This is achieved in two ways.

 

a) Sacrificial anodic protection

 

Here the metal to be protected is made cathode by connecting it to a more active metal (anodic metal) called sacrificial anode. Only the more active metal will

 

a) Sacrificial anodic protection

 

Here the metal to be protected is made cathode by connecting it to a more active metal (anodic metal) called sacrificial anode. Only the more active metal will be corroded, protecting the parent metal. Since the anodic metal is sacrificed, the method is called sacrificial anodic protection. Mg, Zn are used as sacrificial anodes. Metal to be protected

 


Applications

       Protection of buried pipelines, cables

          Protection of ships and boats

            Calcium metal is used to minimize engine corrosion

 

Magnesium sheets are inserted into domestic water boilers to prevent rust formation.

 

b) Impressed current cathodic protection method

 

Here an impressed current is applied in an opposite direction to annul the corrosion current. Thus the corroding metal is converted to cathode from anode.The negative terminal of battery is connected to the metal to be protected. The positive terminal is connected to an inert electrode like graphite. The anode is buried in a „back-fill‟ (containing a mixture of gypsum, coke breeze and sodium sulphate) to increase electrical contact.


Mg Metal to be protected

 

 

 

 

APPLICATION

 

       Protection of tanks, transmission line towers, underground water pipes, oil pipe line, ships etc.

 

Limitations

       It is costly

          It fails when current is switched off.

 

Corrosion inhibitors

 

A corrosion inhibitor is a substance that reduces corrosion, when added to the corrosive environment. There are three types of inhibitors.

 

       Anodic inhibitors - chromate, nitrate

          Cathodic inhibitors - amines

 

            Vapour phase inhibitors - benzonitrile.

 

       Anodic inhibitors

(e.g.) chromate, nitrate, phosphates, tungstate.

The inhibitors form insoluble compound with the newly produced metal ions and prevent corrosion. This compound is adsorbed on the metal surface to form a passive film. Anodic inhibitors are used to repair

 

i) the crack of oxide film on metal surface

          pitting corrosion

            porous oxide film on metal surface

 

          Cathodic inhibitors

There are two types depending on the nature of cathodic reaction in an electrochemical reaction.

 

a) In acidic solution

 

Example : amines, thiourea, mercaptans act as inhibitors. Here evolution of H2 is the cathodic reaction.

 

2 H+ + 2e- → H2

 

The corrosion is controlled by slowing down the diffusion of H+ ions to cathode by addition of the inhibitor which is adsorbed on the metal surface.

 

b) In neutral solution

 

Example : hydrazine, sodium sulphite act as inhibitors. Here OH- ions are formed at cathode. H2O + ½ O2 + 2e- → 2 OH

 

Corrosion is controlled by eliminating O2 from the corroding medium by adding Na2SO3. The OH- ions can be eliminated by salts of Mg, Zn etc.

iii) Vapour phase inhibitors

 

(e.g.) benzotriazole, dicyclohexyl ammonium chromate act as inhibitors. These organic inhibitors readily vapourise and form a protective layer on the metal surface.

 

4.2 Control of corrosion by modifying the environment:

There are five methods

 

1. Deaeration :

 

Presence of oxygen increases corrosion rate. Deaeration involves removal of dissolved oxygen by increasing the temperature together with the mechanicalagitation. This also removes dissolved oxygen.

 

2. Deactivation:

It is the removal of dissolved oxygen by adding chemicals in aqueous solution.

 

(E.g.) 2Na2SO3 + O2 → 2Na2SO4

 

3. Dehumidification:

 

It is the removal of moisture from the air by reducing the relative humidity of the surrounding air. It can be achieved by adding silica gel or alumina which absorbs

 

moisture.

 

4. Alkaline neutralization:

 

The acidic nature of the corrosive environment is due to the presence of HCl, SO2, CO2 etc. They are neutralized with alkali spray. E.g. NaOH lime etc.

5. Using corrosion inhibitors :

 

A corrosion inhibitor is a substance that reduces the corrosion of a metal when added to corrosive environment

 

Applications

To prevent corrosion in closed space, storage containers, sophisticated equipment etc.

 

PROTECTIVE COATINGS

 

Metal surface is covered by a protective coating to prevent corrosion. The coating acts as a physical barrier between the metal surface and the environment. The coating gives a decorative appeal and also imparts hardness, oxidation resistance and thermal insulation to the surface. The main types of coating are:

 

         Metallic coating

         Chemical conversion coating

         Organic coating

         Non-metallic coating

 

 

5 PAINT

 

Paint is a mechanical dispersion of one or more fine pigments in a medium (thinner + vehicle). When a paint is applied to metal surface, the thinner evaporates. The vehicle undergoes slow oxidation to form a pigmented film.

 

5.1 Requirements or requisites of a good paint

A good paint should,

 

       have good covering power

          spread easily on the surface

            not crack on drying

           adhere well to the surface

         give a glossy film

           be corrosion and water resistant

            have stable colour

 

5.2Constituents of Paint

 

Pigment Vehicle Thinner Drier Filler Plasticizer

 

Anti skinning agent

 

1. Pigment

It is a solid that gives colour to the paint.

 

Functions:

1. To give colour and opacity to the film.

 

2. To provide strength to the film.

3. To protect film by reflecting U.V. rays.

4. To provide resistance to abrasion and weather.

 

Example:

White pigment - White lead, TiO2

 

Blue pigment - Prussion blue

Green pigment - Chromium oxide

Red pigment - Red lead, Fe3O4

 

2. Vehicle (or) Drying Oil

 

It is the film-forming liquid. It holds the ingredients of the paint. It is a nonvolatile high molecular weight fatty acid of vegetable or animal.

 

Function

1. To hold the pigment on the surface.

 

2. To form a protective layer by oxidation and polymerization. 3.  To impart water repellency, toughness and durability of film. 4. To improve adhesion of film.

 

Example

Lin seed oil, Castor oil.

 

3. Thinner

 

It is the volatile portion of paint. It is added to reduce the viscosity of the paintfor easy application on the surface. It easily evaporates after paint is applied.

 

Functions

1. To reduce viscosity of paint.

 

2. To dissolve vehicle and other additives.

3. To suspend the pigments.

4. To increase elasticity of film.

5. To increase penetration of vehicle.

6. To improve drying of film.

 

Example

Turpentine, Dipentine, Xylol.

 

4. Drier

It is a substance used to speed up drying of the paint.

 

Functions

1. To act as oxygen carrier or catalys.

 

2.To provide oxygen essential for oxidation and polymerization of drying oil.

 

Example

Metallic soap, linoleate and resinate of Co, Mn etc.

 

5. Extender or Filler

These are white pigments that form bulk of the paint.

 

Functions

1. To reduce cost of paint

 

2. To prevent shrinkage and cracking of film

3. To modify shades of pigment

4. To retard settling of pigments in paint.

 

Example

Talc gypsum, china-day.

 

6. Plasticizer

It is added to the paint to provide elasticity to the film and prevent its cracking.

 

Example

Triphenyl phosphate, Tricresyl phosphate

 

7. Antiskinning agent

It is a chemical added to the paint to prevent gelling and peeling of the paint.

 

Example

Polyhydroxy phenols.

 

Pigment Volume Concentration (P.V.C.)

 

The P.V.C. of a paint is calculated using the equation. P.V.C. =

 

Volume of pigment in the paint =Volume of pigment in the paint + Volume of non-volatile vehicle in the paint

 

If P.V.C. is high, durability, adhesion and consistency of the paint will be low.

 

Failure of Paints

A paint may fail due to any one of the following reasons:

 

Chalking : It is the gradual powdering of the paint film on the painted surface. This happens due to improper dispersion of pigment in vehicle.

 

          Cracking : A paint film cracks due to unequal expansion or contraction of paint coats.

Evasion : This is very quick chalking.

iv) Blistering : It is due to improper surface exposure of paint to strong sunshine.

 

 

6 Metallic Coating

 

6.1 Electroplating or Electro-deposition

 

It is the deposition of coat metal on the base metal by passing direct current through an electrolytic solution of a soluble salt of the coat metal. The base metal to be electroplated acts as cathode. The coat metal or an inert electrode forms anode. The electrolyte is a soluble salt of coat metal.

Objectives or uses or applications of Electroplating:

       To enhance resistance to corrosion of base metals.

 

          To give a decorative appearance.

            To enhance resistance to chemical attack.

           To improve hardness and wearing resistance.

         To obtain polished surface.

 

Theory

 

If the coating metal itself forms the anode, the concentration of electrolyte bath does not change during electrolysis. The metal ions deposited on the cathode are replenished continuously by dissolution of the anode.

 

Example

Electroplating of Gold


The object to be gold plated is treated with organic solvent like acetone, CCl4

 

to remove grease, oil etc. It is then washed with dil H2SO4 to remove scales, oxides etc. The cleaned object is made cathode of electrolytic cell. Anode is a gold plate. AuCl3 solution is the electrolyte. When current is passed into the solution, gold ions migrate to the cathode, get reduced and deposit on the object.

 

 To achieve a strong adherent and smooth deposit, glue or gelatin is added to the electrolyte bath. To enhance the brightness of the deposit, brightening agents are added to the bath.

 

Conditions

       Temperature : 600C

          Current density : 1 to 10 mA/cm2

            Low metal ion concentration

           Buffer solution to maintain pH.

 

Applications or Uses or Objectives

       To give a decorative appearance.

          Electrical and electronic applications

            To get a thin coating of gold on cheap jewellery

           To achieve oxidation resistance, corrosion resistance etc.

 

6.2 Electroless Plating

 

It is the deposition of a noble metal (from its salt solution) on a catalytically active metal surface using a reducing agent without use of electric current.The reducing agent reduces the metal ions. The metal atoms get deposited over the surface to give a thin uniform coating.

 

Metal ions + reducing agent   metal (deposited) + oxidation product

 

Example

Electroless nickel plating

The various steps are:

 

Step I : Pretreatment and activation of the surface:

 

The surface to be plated is degreased by using organic solvents or alkali and then accompanied by acid treatment.

 

       The surface of stainless steel is activated by dipping in hot solution of 50% H2SO4.

          Mg alloy surface is activated by giving a thin coating of zinc and copper overit.

            Al, Cu, Fe, brass etc, do not require activation.

           Plastic, glass etc, are activated by dipping in a solution of SnCl2/HCl and then

in PdCl2 solution. On drying a thin layer of palladium is formed on the surface.

 

Step II : Preparation of plating bath:

The plating bath consists of:

 

       Coating Metal : A solution of NiCl2 20g/lit.

          Reducing agent : Sodium hypophosphite 20g/lit.

 

        Exaltant to accelerate coating rate and complexing agent : Sodium succinate 15g/lit.

 

           Buffer to maintain pH at 4.5 : Sodium acetate 10g/lit.

         Temperature 93oC

 

The pretreated object is immersed in the plating bath for required time. The following reactions occur and nickel is coated on the object.

 

Cathode : Ni2+ + 2e- → Ni

Anode : H2PO2- + H2O → H2PO2- + 2H+ + 2e

Overall Reaction : Ni2+ + H2PO2- + H2O → Ni + H2PO3- + 2H+

 

Uses of Nickel Plating

       For decorative coating of jewellery, decorative items and automobile spares

 

          For coating of polymers for decorative purpose.

            For electronic appliances.

 

Advantages of electroless plating over electro plating

       Electricity is not necessary

          Complicated parts are uniformly coated

            Plastics, glass etc, are easily coated

           Good mechanical, chemical and magnetic properties are obtained.

 

Differences between Electroplating and Electroless plating:

Electroplating Electroless Plating

        It is done by passing current. It is done by auto catalytic redox reaction.

        Separate anode is required. Catalytic surface of the object acts as anode.

        Anode reaction: M → Mn+ + ne- Anode reaction : R → O + ne-

        Cathode reaction: Mn+ + ne-→  M Cathode reaction: Mn+ + ne- → M

        Irregular objects are not satisfactorily plated All objects are satisfactorily plated.

        Object to be coated forms the cathode .Object to be coated forms catalytically active surface.

        It is carried out on conducting materials.It is carried out even on insulators.

 

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