CHAPTER-07

ELECTROCHEMISTRY

CHEMISTRY SSC-I                                 CH:07

ELECTROCHEMISTRY                                            LEC# 01

Q: What is electrochemistry and how it is important in every day life?

Ans: ELECTROCHEMISTRY:The branch of chemistry which deals with the conversion of electrical energy into chemical energy and chemical energy into electrical energy is called ELECTROCHEMISTRY.

Importance in everyday life

  • Electrochemical Industries produce millions of tons of Important metals such as Copper, Aluminium, Magnesium, Sodium and Zinc etc. through electrolysis processes.
  • They also produces Caustic Soda, Chlorine, Silicon carbide etc.
  • Important processes such as Rusting of Iron objects, combustion of fuel  in automobile engine, forest fire and metabolism of food in human and animal bodies are involving Redox reactions.
  • The turning on of a flash light, mobile, calculator, electronic toys etc also generates out because of a current of electricity.
  • House hold bleaching agents decolourize the color bearing substances in strains involve Redox reactions.
  • Dying of Aluminium can produce metallic red, blue or other colours on the metal surface because of process Electroplating.

Q: How Redox or oxidation reductions reactions can take place              

Ans: Redox Reactions

There are three different ways of expressions of oxidation and Reduction reactions.

  • Loss or gain of oxygen and hydrogen.
  • Loss or gain of electron (s).
  • Oxidation number decreases or Increasing.

Importance of Redox reactions:

Redox is stand for Reduction-Oxidation reactions simultaneously observed in same chemical reaction. These reactions play an important role as follows:

  • In Steel mills, iron ores usually oxides of Iron are converted to the pure metal on Industrial level by the following reaction in the blast furnace

C + O2                    CO2 + heat (Exothermic reaction)

CO2 +C                   2CO(Endothermicreaction)

Fe2O3 + 3CO                                          2Fe + 3CO2

 In this reaction, Fe2O3 loses oxygen which is gains by CO and oxidizes into CO2. Thus Fe2O3 is reduced into Fe and CO is oxidized into CO2.

  • Acetylene (C2H2) is commercially used for cutting and welding metals. When acetylene burns, it produces a very hot flame known as oxy-acetylene flame and the following reaction takes place.

2C2H2 + 5 O2                          4CO2+H2O + heat

In this reaction, Acetylene loses hydrogen and oxygen gains hydrogen to form water. Thus Acetylene gets oxidized (loses hydrogen) into CO2 and oxygen is reduced into H2O (gain Hydrogen)

  • Coal is burned in Thermal power stations to produce electricity. The following reaction occurs, when it burns  C + O2        CO2 + heat

In this reaction, carbon atom gains oxygen and  gets oxidized into CO2.

  • The formation of a loose flaky layer of  Iron (III) oxide, Fe2O3 on the surface of Iron is called Rust. In rusting  of Iron, following reaction occurs4Fe + 3O2                            2Fe2O3
Text Box: SELF ASSESSMENT EXERCISE 7.1

In this reaction, iron gains oxygen and get oxidized into Fe2O3.

Identify elements undergoing oxidation and reduction in the following reactions:

Q: Define oxidation and support your answer with examples?

Ans: Oxidationis defined as the gain of oxygen atoms

i) Oxidation of C into CO2 by combustion

*
  • C + O2                                     CO2
*
  • 4Al+3O2                 2Al2O3

Oxidation is also defined as the loss of hydrogen atoms by a molecule.

2C2H2+5O2                             4CO2+2H2O + heat

CH4 + 2O2                               CO2 + 2H2O + heat

A process that involves the loss of electrons by an element is called oxidation. Metals lose electron and form M+.

  • M                                            M+1 + e ̅
  • Li                                             Li+1 + e ̅  
  • Mg                                          Mg+2 +2 e ̅
  • Na                                           Na+1 + e ̅ 
  • Ca                                            Ca+2 +2e

CHEMISTRY SSC-I                                 CH:07

ELECTROCHEMISTRY                                            LEC# 02

Q: Define reduction and also give suitable examples?

Reduction: is defined as the gain of Hydrogen atoms.

  • 2H2 + O2                                 2H2O
  • N2 + 3H2                                 2NH3

Reduction is also defined as the loss of O-atoms.

  • Fe2O3 + 3CO                                          2Fe + 3CO2
  • Zn + CuO                                                ZnO + Cu

A process that involves the gain of electron (s) by a substance is called Reduction.

  • N + 3e                                    N-3 (Nitride)           
  • P + 3 e                                    P-3 (Phosphide)
  • O + 2 e                                   O-2 (Oxides)         
  • S + 2e                                      S-2 (Sulphides)

Q: Define oxidation state and give suitable examples?

Oxidation Number or Oxidation State

It is defined as the number of charges on an atoms, molecule or compound is called Oxidation state.

In Na+1  the oxidation number of Na is +1

•              In O-2, the oxidation number of O is -2

•              In Clatom, the oxidation Number is Zero because it is a neutral atom.

Q: Compare oxidation and reduction

Oxidation Reduction
Gain of oxygen Loss of oxygen
Loss of hydrogen Gain of hydrogen
Loss of electrons Gain of electrons
Increase in oxidation number/state Decrease in oxidation number/state
Text Box: SELF ASSESSMENT EXERCISE 7.2

In the following reactions, identify which element is oxidized and which element is reduced.

Q: What rules are followed to assign an oxidation number?

Rules for Assigning Oxidation states or Numbers

  • The oxidation state of any uncombined or free elements is always zero e.g., oxidation state of Zn, Na, H in H2, S in S8 etc is zero.
  • In simple ions, oxidation state is same as their charge e.g., oxidation state of Na is Na+1 and Ca is Ca+2 are +1 and +2 respectively.
  • In a complex ion the total sum of oxidation states of atoms is equal to the charge on their ion. e.g., in CO32-, the sum of oxidation states of C and 3 O atoms is -2. Similarly, in NH4+1, the sum of oxidation states of N and 4H atoms is +1.
  • The oxidation number of each of the atoms in a molecule or compound counts separately and their algebraic sum is zero e.g., In HCl, the sum of oxidation states of H and Cl atoms is zero. Similarly in CO2, the sum of oxidation states of one C and 2 oxygen atoms is zero.
Elements Oxidation State
Group-IA +1
Group-IIA +2
Group-IIIA +3
H +1 (except in metal hydrides where it is -1)
Group-VIIA -1
O -2(except peroxides and in OF2)

FIND OXIDATION STATE OF FOLLOWING

Cr in K2Cr2O7 and  B in H3BO3

K2Cr2O7 =0 2K +2Cr + 7O = 0 2(+1)+ 2Cr + 7(-2)=0 2+ 2Cr -14 =0 2Cr – 12= 0 2Cr = 12 Cr = 12/2 = +6 So O.N of Cr is +6   H3BO3 =0 3H +B + 3O = 0 3(+1)+ B + 3(-2)=0 3+ B -6 =0 B-3= 0 B = +3   So O.N of B is +3
Text Box: SELF ASSESSMENT EXERCISE 7.3

Find Oxidation Number of S in H2SO4, N in HNO3 and N in NO2.

What are oxidizing and reducing agents?

Ans:

  • Oxidizing Agent

An oxidizing agent is the reactant containing the element that is reduced (gains electrons) in a chemical reaction.

  • Reducing Agent

A Reducing agent is the reactant containing the element that is oxidized (loses electrons) in a chemical reaction.

For Example

Sodium and chlorine reacts to form Sodium chloride.

2Na + Cl2                                                2NaCl

Na is reducing agent because it is being oxidized as its oxidation no increases from

0à+1

2Na° + Cl2°                                             2Na+1 Cl-1

while Cl2 is oxidizing agent because it is being reduced as itsoxidation no. 

decreases from 0 à -1

CHEMISTRY SSC-I                                 CH:07

ELECTROCHEMISTRY                                            LEC# 03

Q: How can you identify oxidizing and reducing agents in a chemical reaction?

Ans:

Consider the following reaction that takes place in the manufacture of steel.

Fe2O3 + 3CO                                          2Fe + 3CO2

To identify the oxidizing and reducing agents, work out the oxidation states of all the elements involved in the reaction.

+3(-2)   +2 -2                                          0 +4(-2)2

Fe2O3  +   3C O                                      2Fe +3CO2

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  • Carbon is being oxidized because there is an increase in its oxidation state from +2 to + 4.
*
  • Fe is being reduced because there is a decrease in its oxidation state from +3 to zero.
*
  • The reactant CO contains the C that is being oxidized, so CO is reducing agent.
*
  • The reactant Fe2O3 contains the Fe that is being reduced. So Fe2O3 is oxidizing agent.
Text Box: SELF ASSESSMENT EXERCISE 7.4

Zn + 2MnO2 + H2O                               Zn(OH)2 + Mn2O3

Identify the oxidizing and reducing agents in this reaction.

      2.       Identify oxidizing and reducing agents in the following reactions:

a)   2S+Cl­­­2                              S2Cl2

Q: What are redox reactions?

AnS:

Redox Reactions

Chemical reactions in which the oxidation state of one or more substances changes are called oxidation-reductionor redox reactions.

  • Cations are formed from metallic elements after removal of electron (s)
  • Anions are formed from non-metallic elements after gaining of electron (s)

as a result of combining of cations and Anions, Ionic compound are formed e.g: Calcium (metal) and Chlorine (non-metal) contain neutral atoms. After transference of elements from calcium atoms to chlorine atoms and form calcium chloride involved oxidation- reduction reactions simultaneous called REDOX Reaction (Red for Reduction and Ox for Oxidation)

0                                              0                                                    +2   2(-1)

Ca            +              Cl2 —————–àCaCl2

oxidation                                 reduction

Q: How redox reactions are used in metallurgical processes?

Method of Recovering Metals from their Ores

Redox Reactions are commercially very important. Most of the metals are recovered from their ores by Redox Reactions

  • Most of the metals are found in nature as oxides or Sulphide ores. After mining the ore, desired mineral is separated from the other materials.
  • Purified Metal Oxides are reduced to free metals by using Reducing Agent. Aluminium, Coke, Carbon, monoxide gas and Hydrogen gas are generally used as Reducing Agents.
  • Iron can be extracted from their ore hematite (Fe2O3) by Coke (carbon) used as a reducing agent in a Blast Furnace.
  • Iron ore (Fe2O3), Lime stone (CaCO3) and Coke are introduced into the Blast Furnace from the top. A blast of hot air is forced up the furnace from the bottom so called Blast Furnace. The oxygen gas reacts with Coke to form mostly carbon monoxide and some carbon dioxide. These reactions are highly exothermic. As hot CO rises, it reacts with Iron oxide and reduces it to iron. Molten Iron collects at the bottom of the furnace. Limestone removes impurities i,e. SiO2, Al2O3 Iron as slag. The reactions are given below:

C + O2                                    CO2

CO2 + C                                   2CO

3CO + Fe2O3                          2Fe + 3CO2

CaCO3                                     CaO + CO2

CaO + SiO2                             CaSiO3

CaO + Al2O3                           Ca( AlO2 )2

Similarly, Lead and Zinc metals are extracted by their ores.      

  • ZnO         + C                           Zn + CO
  • PbO + CO                               Pb + CO2

                CuO  + H2                               Cu + H2O

Q: What are electrochemical cells? , also describe nature of chemical processes occur in these reactions?

Electrochemical Cells

An arrangement which consists of electrodes dipped into an electrolyte in which Chemical reaction uses or generates electric current is called Electrochemical Cell, e.g:

  • Electrolytic Cells (Down’s and Nelson’s cell)
  • Galvanic cells (Daniel’s and Ni-Cd cell)

Electrolytic Cell

The cell in which a reaction occurs with the help of Electric current is called Electrolytic cell.

  • Down’s Cell           2NaCl——à2Na(s) + Cl2(g)
  • Nelson’s Cell

2NaCl(eq)+ 2H2O——–à2Na++ 2OH + H2 + Cl2

  • Non-spontaneous oxidation- reduction reactions takes place
*
  • Galvanic or Voltaic Cell

The cell in which a reaction generates electric current is called voltaic or Galvaniccells.e.g.

  • Daniel’s cell, Lead storage battery, fuel cell etc.
  • Spontaneous oxidation-reduction reaction takes place.
  • CHEMISTRY SSC-I                                 CH:07

ELECTROCHEMISTRY                                            LEC# 04

Nature of Electro Chemical process

Electrochemical processes are oxidation-reduction reactions in which chemical energy released by a spontaneous reaction is converted to electricity or in which electrical energy is used to drive a non-spontaneous reaction. Whether an electrochemical process releases or requires energy, it always involves the transfer of electrons from one substance to another. This means that this process always involves an oxidation- reduction or a redox reaction.

Q: What are electrolytes & non electrolytes? Also differentiate between weak and strong electrolytes ?

ElectrolytesThe substances which undergoes partial or complete dissociation into ions in solution or molten form passing electricity are called Electrolytes.

e.g.: NaCl solution, KCl solution, HCl, NaOH etc.

Non-Electrolytes

A substances that cannot conduct electricity when dissolved in water or in the molten state is called Non-electrolytes.

e.g: urea, glucose, Sucrose, benzene ,woodetc.

Q: Differentiate b/w spontaneous and non-spontaneous process?

Ans:

  • Spontaneous Process

A physical or chemical change that occurs by itself is called a spontaneous process.

Spontaneous process do not require a source of energy to make them happen. For instance water flows from higher level to lower level. Iron placed in moist air, rusts. The flow of electrons through a conductor can be obtained from a spontaneous oxidation-reduction reaction. This is the basis for how batteries work.

  • Non-spontaneous Process

Aa physical or chemical change that requires a source of energy tomake them happen is called non-spontaneous process.

For example water can be made to flow from lower level to higher level by using a pump.

>Purification of copper by elecrolysis

 Q: Sketch a electrolytic Cell, the Cathode, Anode and the direction of flow of the electrons.?

http://upload.wikimedia.org/wikibooks/en/b/b4/Chemical_Principles_Fig_1.9.png

Ans:

An electrolytic cell consists of

*
*
*

The figure shows that electrons move from anode to cathode in the outer circuit, in the solution the cations move towards cathode and anions towards anode. At anode anions oxidize by loosing electrons. At cathode cations reduce by gaining electrons. This means oxidation occurs at anode and reduction at cathode.

At Anode                                

At cathode                             

Q: List some importants uses of electrolytic cells?

Uses of Electrolytic Cells

  • Down’s Cell is used for the commercial preparation of sodium metal. It produces chlorine gas as by product.
  • Nelson’s Cell is used for the commercial preparation of sodium hydroxide. It also produces chlorine and hydrogen gas as by product.
  • Electrolysis cells are used for the commercial preparation of calcium and magnesium metals.
  • It is used to produce aluminium metal commercially.
  • It is used for the purification of copper.
  • Electrolytic cells are used to electroplate metals such as tin, silver, nickel etc on steel.
  • Electrolytic cells are used to prepare anodized aluminium. Anodized aluminium can absorb dyes.
  • Dyeing of anodized aluminium can produce metallic red, metallic blue or other metallic colours on the metal surface.                                                                        

Q: What is a Galvanic cell/Daniel Cell? Also sketch and label it?

AnS:

Galvanic (Daniel) Cells.                         

The cell which involves spontaneous redox reaction to generates electricity is called a galvanic or voltaic cell. The name Voltaic is given to this cell because Alessandro Volta discovered first such cell. The English chemist Fredrick Daniel Constructed first Voltaic cell using zinc (Zn) and copper (Cu) electrodes. Therefore this cells is named as Daniel Cell or  galvanic cell

  • A zinc bar dipped into a 1 M ZnSO4 solution.
  • A copper bar dipped into a 1 M CuSO4 solution.
http://imbuepreachers.com/wp-content/uploads/2013/04/fetch.png
  • A salt bridge which is inverted U tube containing an inert electrolyte such as KC1. Its ions do not react with the electrodes or with the ions in solution. It makes the electrical contacts between the solutions through which ions can move.
  • A voltmeter to measure current.

Each compartment of cell is called a half cell. Thus a Daniel cell consists of two half cell joined in series. When circuit is complete electrons begin to flow from Zn rod through the external wires to Cu rod. Thus Zn half cell acts as anode and Cu half cell as cathode.

  • Note that a half cell consists of a metal rod dipped in the solution of its salt.

Reactions in a Daniel Cell

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  1. In Daniel cell, electrons flow from Zn rod, through the external wire to Cu rod. This is because Zn has more tendency to undergo oxidation than Cu. Zn atoms from the rod go into the solution as Zn+2 ions leaving electrons on the rod. These electrons flow in the external circuit. Thus oxidation half reaction occurs at anode compartment.
*
  • Cu+2 ions in copper sulphate solution capture electrons from Cu electrode and are reduced. Reduction half cell reduction occurs at the cathode compartment. Such oxidation and reduction reactions are called half cell reactions.

At Anode(Oxidation half reaction):

At Cathode (Reduction half reaction):

Q: What is salt bridge? What is function of using salt bridge in Galvanic cell?

Ans:

Salt Bridge

It is a U-shaped glass tube having a Saturated Solution of some strong electrolyte like KCl, K2SO4 or KNO3. The glass tube is sealed at both of its ends by a porous material like glass wool or cotton play. It prevents the physical contact between the two electrolytic solutions.

Function of Salt Bridge           .

  • It connects the solutions in two half cells and completes the cell circuit.
  • It maintains the electrical neutrality by the diffusion of ions through it.

CHEMISTRY SSC-I                                                 CH:07

ELECTROCHEMISTRY                                            LEC# 05

Q: Define cell and battery?

Cell and Battery

  • A cell is a combination of two metal plates, one acts as cathode (negative terminal) and other act as Anode (positive terminal). It is represented as +|| and its potential is 2 volt.

Conventionally, Anode line is bigger than cathode.

  • A Battery is a Galvanic cell groups joined in series of a circuit. It generates electric current by a Redox Reaction. It is represented as +||||||

Examples of batteries are dry cell, storage cell, mercury battery, Ni-Cd batteiyetc,,

https://encrypted-tbn1.gstatic.com/images?q=tbn:ANd9GcTlMA7xV9kA6w53GAh8bHXHRB8l29DZBRl3IG938ivjAnj-lCOi

Q: How a battery produces electrical energy in a dry cell?

Dry CellThe dry cell batteries are used to power many flashlights, toys and small appliances.

  • The anode is the zinc metal of the container and the cathode is an inert graphite rod at the center of the container in contact with a mixture of MnO2 and carbon (charcoal).
  • The electrolyte is a mixture of moist NH4CI and ZnCl2. Following reactions take place in it.

At Anode (oxidation takes place which means loss of electrons)

At Cathode (Reduction takes place which means gain of electrons)

2NH4+ + 2MnO2 +2e                         Mn2O3 + 2NH3 + H2O

  • This cell produces a potential of 1.5 V.

ELECTROCHEMICAL INDUSTRIES

Q:Describe how sodium metal is obtained from fused sodium chloride by Down’s cell?

Manufacture of Sodium Metal from Fused Sodium Chloride

On the large scale sodium metal is produced by the electrolysis of fused NaCl. An electrolytic cell called Down’s Cell is used for this purpose as shown below.

http://www.sciencemadness.org/scipics/downs_sodium_productioncell.jpg

The electrodes are iron cathode and graphite anode. Chlorine is obtained as by-product. In molten sodium chloride (NaCl), Na+ ions are free to move about. Under the influence of electric current, Na+ ions move towards the cathode and CP ions towards the anode.

Molten sodium is collected into a sodium collecting ring,from where it is periodicallydrained. Whereas, chlorine gas in collected into the funnel at the top of the cell.

Molten NaCl ionizes as:
Reaction at anode (oxidation)
  Reaction at cathode (reduction)
  Overall sum of reactions
 

Q: Describe how Sodium Hydroxide is prepared from Brine solution in Nelson cell?

http://www.newconceptinfosys.com/ibm/imgs/new/nelson.gif

Manufacture of Sodium Hydroxide from Brine

Electrolysis of brine, a concentrated aqueous solution of sodium chloride is used for the industrial production of sodium hydroxide. Electrolysis of brine produces

simultaneously three important industrial chemicals, chlorine gas, hydrogen gas and sodium hydroxide. The electrolytic cell called Nelson’s Cell

Working of Nelson’s Cell

Aqueous solution of sodium chloride consists of Na+, CI, H+ and OH ions. These ions move towards their respective electrodes and redox reactions take place at these electrodes. When electrolysis takes place Clions are discharged at anode and CI2 gas rises into the dome at the top of the cell. The H+ ions are discharged at cathode and H2

gas escapes through a pipe. The sodium hydroxide solution slowly percolates into a catch basin.

  Brine ionizes to produce ions:
 
Reaction at anode (oxidation):
 
Reaction at cathode (reduction):
 
Overall cell reaction of this process:
 

Q: Explain electrolytic refining of copper?

Ans:

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Electrolytic Refining of Copper

Electrolytic cell can be used to purify copper. For this purpose, an impure copper is made the anode and a thin sheet of pure copper is made a cathode. The electrolyte used is copper sulphate (CuS04) solution.

The impurities are mostly silver (Ag), gold (Au), platinum (Pt), iron (Fe) and zinc (Zn) and they settle down as anode mud. Presence of these metals decreases the electrical conductivity of copper (Cu).Cathode is a very thin sheet of very pure copper. When the cell is operated at the given value, the less active metals simply fall of the electrode and settle to the bottom of container. At cathode, Cu+2is reduced but Zn+2 and Fe+2 remain in the solution.

In this way, 99.5% pure copper is obtained

CHEMISTRY SSC-I                                                 CH:07

ELECTROCHEMISTRY                                            LEC# 06

Q: Define electroplating? What is its purpose? What are conditions for good electroplating?

Electroplating It is the art of depositing one metal over the surface of another metal with the help of Electric Current.

Purpose of Electroplating

The electroplating is done for following purposes.

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  • Preservation from rusting and enhancing the life of metals.
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  • Decoration
*
  • Repairing

Conditions for Good Electroplating

  1. High current density
  2. Low temperature
  3. High concentrate ion of metal in its electrolyte.

Q: Describe how various metals are deposited  on steel?

Electroplating on steel

Steel objects are often protected from corrosion by electroplating with zinc, tin and chromium.

  1. Zinc plating

Zinc plating on steel is done by using zinc metal as anode. A solution of potassium zinc cyanide K2[Zn(CN)4] containing little sodium cyanide. The steel object is made cathode. During the electrolysis zinc at the anode enters the solution as Zn+2 ions, which are deposited at the cathode. The electrolyte ionizes as follows.

Following reactions occur at the electrodes:

At anode:               

At cathode:                            

Sodium cyanide prevents the hydrolysis of the electrolyte.

Effect of cyanide on environment

***Cyanide ions are extremely toxic. Therefore, solution containing cyanide ions must not be dropped into rivers and streams. This is responsible for killing fish and other animals.

  • Tin platingFood cans are generally tin plated. Tin plating on steel is done by using tin as anode and a solution of stannous sulphate, (SnSO4) as electrolyte. Few drops of dil H2SO4 are added in the electrolyte to prevent its hydrolysis. The electrolyte ionizes as follows.

                During the electrolysis following reactions occur:

At anode:                               

At cathode:                                            

  • Chromium plating

                Since chromium metal does not adhere strongly to the steel therefore steel is first plated with copper or nickel and then chromium. For electroplating chromium, chromium metal is used as anode and chromium sulphate, Cr2(SO4)3 as an electrolyte. A few drops of dil H2SO4 are added in the electrolyte to prevent its hydrolysis. The electrolyte ionizes as follows:

At anode:

At cathode:             Chromium plated steel is used to make automobile parts

Q: What is meant by corrosion?

Corrosion and its prevention

Corrosion is the process in which a metal reacts with oxygen and moisture in the atmosphere. It is a natural process that converts refined metals to the more stable metal oxides.

Q: How rusting of iron take place?/ describe what chemical reactions are involved in it?

Ans:

Rusting of iron

                Most familiar example of corrosion is the formation of rust on iron. Oxygen and water are necessary for iron to rust. A region of metal surface that has relatively less moisture, acts as anode.

Another region on the surface of metal that has relatively more moisture acts as cathode. The electrons released in the oxidation process reduce atmospheric oxygen to hydroxyl ions.

The Fe+2 ions formed at the anodic regions flow to the cathodic regions through the moisture on the surface. Here Fe+2 ions further react with oxygen to form rust, Fe2O3.xH2O

Q: how corrosion in aluminum take place?

Ans:

Corrosion of Aluminum

Corrosion is not limited to iron. Aluminium is extensively used in the construction of aircraft, ships, cars, cooking utensil, window frames, soda canes etc. Aluminium has much higher tendency to oxidize than iron. This is because a tough layer of insoluble aluminium oxide (Al2O3) forms on its surface when metal is exposed to air. This layer firmly adheres to the metal and serves to protect the underlying aluminum layers from further corrosion.

Q: how we can prevent corrosion in metals?

Prevention of Corrosion

              Prevention of corrosion is an important way of conserving our natural resources. Following methods have been devised to protect metals from corrosion:

  1. Coating with paint:

Corrosion can be prevented by painting the metal, so that it does not come in contact with oxygen and moisture and other harmful agents.  Paint is cheap and easily applied. Paint is used to protect many everyday steel objects such as cars, trucks, trains, bikes, bridges etc. Paint also provides visual appeal.

  • Alloying:
    The tendency of iron to oxidize can be greatly reduced by alloying it with other metals. For example, stainless steel is an alloy of iron chromium and nickel. It is protected from corrosion by an outer layer of Cr2O3.
  • Coating with a thin layer of another metal:
    Metals that readily corrode can be protected by coating with a thin layer of another metal that resists corrosion. This can be done by:

(a) Tinning              (b) Galvanizing (c) Electroplating

  1. Tinning:In the process of tin plating, clean iron sheet is dipped in a bath of molten tin. It is then passed through hot pair of rollers. Tin protects iron effectively, since, it is very stable.
  2. Galvanizing (Coating with Zinc):The process of galvanizing consists of dipping a clean iron sheet in a hot zinc chloride bath and heating. After this sheet is rolled into zinc bath and cooled.
  3. Electroplating:In electroplating an electrolytic process is used to deposit one metal on another metal.
  4. Cathodic Protection:

Cathodic protection is the process inwhich the metal that is to be protected from corrosion is made cathode and isconnected to metals such as magnesium or aluminum. These metals are more activethan iron, so they act as anode and iron as cathode. The more active metalsthemselves oxidize and save iron from corrosion. Cathodic protection isemployed to prevent iron and steel structures such as pipes, tanks, oil rigsetc in the moist underground and

EXERCISE

Q.2:        Give short answers

  1. What is oxidation state?

Ans: It is defined as the number of charges on an atoms, molecule or compound is called Oxidation state.

  1. What is the oxidation number of Cr in chromic acid (H2CrO4)?
  2. Identify reducing agent in the following reaction
  1. Write chemical reactions that occur in Nelson’s cell.
  Brine ionizes to produce ions:
 
Reaction at anode (oxidation):
 
Reaction at cathode (reduction):
 
Overall cell reaction of this process:
 
  • Why tin plated steel is used to make food cans?

Ans:        Tin plated steel is used to make cans. Food and beverages industries use tinplated steel cans, This is because the components of food beverages and the preservatives contain organic acids or their salts.  They may form toxic substances by reacting with iron. These acids and salts are corrosive. Tin plating is non-poisonous and prevents corrosion.

  • Explain one example from daily life which involves oxidation-reduction reaction?

Ans: Batteries

Typical AA or AAA battery that powers your flashlight or remote contains another common example of how redox reactions work for you. In your typical alkaline battery, for instance, the oxidation of zinc at the anode produces the electron flow that powers your appliance. The electrons travel, via your appliance, to the cathode where they reduce manganese dioxide in a classic redox reaction.

Q.3:        Define oxidation and reduction in terms of loss or gain of oxygen or hydrogen.

Q.4:        Define oxidation and reduction in terms of loss or gain of electrons.

Q.5:        List the possible uses of electrolytic cell.

Q.6:        Sketch a Daniel Cell, labeling the cathode, anode, and the direction of flow of the               electrons.

Q.7:        Describe how a battery produces electrical energy.

Q.8:        Describe the method of recovering metal from its ores. {See metallurgy of iron }

Q.9:        Explain electrolytic refining of copper.(Lec # 5)

Q.10:      Compare the effects of Al2O3 and Fe2O3 formation on their parent metals and cite examples from daily life.

Q.11:      Explain how food and beverage industries deal with corrosion. [See question V]

Q.12:      Explain how chemistry interacts with photography.

Light exposes tiny crystals of AgBr that are suspended in the thin layer of gelatin coated on the surface of the photographic paper.

AgBr (clear) + light ——–> Ago (Black metal) + e

Exposure of the AgBr with crystals produces particles of several free silver atoms .As light passes through the lightest regions of the negative, dark regions with the highest density of these metallic silver form — the longer the exposure, the greater the number of silver particles. The dark regions of the negative allow less light to come into contact with the AgBr crystals.

Developing . Hydroquinone which is a mild reducing agent is used as developer. It enhances the image. More AgBr is converted to silver metal by the developer. Those grains of AgBr with metallic silver clusters l on their surface from the initial exposure of AgBr with light, react faster with the developer to produce more black silver metal. The rate of reaction is slower when AgBr grains have fewer silver metal clusters on them.

Ag + (in AgBr) + 1e ——-> Ago (Dark image) + reacted developer

The stop bath, acetic acid, stops the developer from reacting with AgBr. The image does not darken after the print is placed in the stop bath.

The fixer, sodium thiosulfate, reacts with the remaining AgBr to form a water soluble salt that can be washed out in the final step in the water bath.

Q.13:      Electrolysis has a major role in electrochemical industries.

(a)Sketch an electrolytic cell, label the anode and cathode and indicate the direction         of electron transfer. (LEC # 04)

(b)Describe the nature of electrochemical process.

( (LEC # 04)

(c)Distinguish between electrolytic and voltaic cell.

Electrolytic Cell

*
  • The cell in which a reaction occurs with the help of Electric current is called Electrolytic cell. e.g Down’s Cell, Nelson’s Cell
*
  • Non-spontaneous oxidation- reduction reactions takes place

Galvanic or Voltaic Cell

*
  • The cell in which a reaction generates electric current is called voltaic or Galvaniccells.e.g. Daniel’s cell, Lead storage battery, fuel cell etc.
*
  • Spontaneous oxidation-reduction reaction takes place.

Q.14:      State the substances which are oxidized or reduced. Give reason for your answer.

Q.15:      (a)           Define oxidation number or oxidation state. (see Q: 2)

(b)           Find the oxidation state of nitrogen in the following compounds.

(i) NO2    (ii)  N2O  (iii)  N2O3                (iv)  HNO3

Q.16:      Find the oxidation state of S in the following compounds.

(a)           H2S          (b)           H2SO3      (c)           Na2S2O3

Q.17:      (a)           Define oxidizing and reducing agents.

                (b)           Identify the oxidizing agents and reducing agents in the following reactions:

Q.18:      Hydrogen peroxide reacts with silver oxide and lead(II) sulphide according to the               following equations.

                Is hydrogen peroxide an oxidizing or reducing agent in these reactions. Give your reason.

Thinktank

1:            What materials do you need to electro plate copper onto an iron nail. Make a diagram showing how these materials should be arranged.

Materials: Nail,strip of copper metal,Copper (II) sulfate

http://img.bhs4.com/9D/3/9D34CA5F585CAAC3B09B912D8E5116B5C53CEC18_lis.jpg

2:            Describe the process that is occurring in the following illustration. Shoe has steel strips.

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3:            Following redox reaction occurs in the voltaic cell illustrated below:

nelson-cell-002.png

                Identify the anode, cathode and indicate the direction of flow of electrons.

Ans: In above cell Fe ionize into Fe+2 , thus undergo oxidation , While Ni+2 accept electron coming from anode and reduces to Ni. So Fe is reducing agent while Ni is oxidizing agent.

Reactions at Anode:

 Fe                                           Fe+2 + 2e

Reaction at cathode:

Ni+2  + 2e                               Ni

Directionof Current is from anode to a

Cathode