ELECTROCHEMISTRY AND APPLICATIONS SYLLABUS Part A · The emf of the cell is 1.1 volts. Direction of...

UNIT III ELECTROCHEMISTRY 2019-20

Engineering Chemistry Page 41

ELECTROCHEMISTRY AND APPLICATIONS SYLLABUS

Part A

Electrochemistry: Electrodes – concepts, electrochemical cell, Nernst equation, cell

potential calculations. Primary cells –dry cell- Secondary cells – lead acid, nickel-cadmium

and lithium ion batteries-working of the batteries including cell reactions- Fuel cells,

hydrogen-oxygen, and methanol fuel cells – working of the cells.

Part B

Corrosion: Introduction to corrosion, mechanism of dry and wet corrosion, Pilling Bedworth

ratios and uses, Types of corrosion – Differential aeration corrosion, galvanic corrosion,

pitting corrosion, waterline corrosion and stress corrosion, Factors affecting the rate of

corrosion – metal based factors and environmental based factors, protection techniques –

metal coatings – galvanization and tinning, cathodic protection, inhibitors – cathodic and

anodic, organic coatings – paints – constituents and their functions.

Course Outcomes: The student will acquire the fundamental knowledge of Electrochemistry/

electrochemical reactions to understand the principles of batteries, fuel cell and corrosion

Part A

OUTLINES

Introduction

Galvanic cells – electrode potentials

Electrochemical series

Electromotive force (Nernst equation)

Batteries

Fuel cells



1. INTRODUCTION

Electric current is the flow of electrons generated by a battery, when the circuit is completed. A

substance which allows electric current to pass through it is called a conductor, e.g., all metals,

graphite, fused salts, aqueous solutions of acids, bases and salts; while insulator or non-conductor is a

substance which does not conduct electric current i.e., which does not allow the passage of electric

current through it.

The conductors are of two types

a. Metallic conductors are the substances which conduct electricity, but are not decomposed by

it, e.g., all metals, graphite etc.

b. Electrolyte is a substance which in aqueous solution or in molten state liberates ions and

allows electric current to pass through, thereby resulting in its chemical decomposition, e.g.,

acids, bases and some salts.

1. Electrochemical Cell (Galvanic cell)

An electrochemical cell is a device in which a redox chemical reaction is utilized to get electrical

energy. An electrochemical cell is generally referred to as voltaic cell or galvanic cell. The

electrode where the oxidation occurs is called anode and the electrode where reduction occurs, is

called cathode. Example: Daniels Cell, Leclanche cell

The Daniel cell (Figure above) consists of zinc electrode dipped in ZnSO4 solution and copper

electrode, dipped in CuSO4 solution. The two solutions are separated by salt bridge so as to avoid

direct contact or mixing with each other. The electrode reactions in Daniel cell are

At anode: Zn Zn2+

+ 2e- (Oxidation)

At Cathode: Cu2+

+ 2e- Cu (Reduction)

Cell Reaction: Zn + Cu2+

Zn2+

+ Cu

Zn has more tendency to form Zn2+

and hence Zn metal acquires a negative charge; and Cu2+

has more

tendency to get deposited as Cu. Hence, copper electrode becomes positively charged. As a result, the



electrons via the external circuit constitutes the electric current in the opposite direction. The emf of

the cell is 1.1 volts.

Representation of a Galvanic Cell

A galvanic cell can be represented as follows;

a) Anode is written on the left hand side; while the cathode is written on the right hand side.

b) The electrode on the left (anode) is represented by writing the metal or solid phase first and the

electrolyte separated by a vertical line or semicolon

Zn(s) | Zn2+

(aq) or Zn(s); Zn2+

c) The cathode of the cell is written on the right hand side. In this case, the electrolyte is

represented first and the metal or solid phase, thereafter separated by a vertical line or semicolon.

Cu2+

| Cu(s) or Cu2+

; Cu(s)

d) A salt bridge is indicated by two vertical lines, separating the two half cells.

Thus, applying above considerations to Daniel Cell, we may represent it as

Zn(s) | Zn2+

(1M) || Cu2+

(1M) | Cu(s)

2. Electrochemical series: It is the series of the elements arranged in increasing order of their

standard electrode reduction potentials.

Metal ion Standard reduction potential in volts

Li + e- Li - 3.05

K+ + e

- K - 2.93

Ca2+

+ 2e- Ca

- 2.90

Na+ + e

- Na - 2.71

Mg2+

+ 2e- Mg - 2.37

Al3+

+ 3e- Al - 1.66

Zn2+

+ 2e- Zn - 0.76

Cr3+

+ 3e- Cr - 0.74

Fe2+

+ 2e- Fe - 0.44

Ni2+

+ 2e- Ni - 0.23

Sn2+

+ 2e- Sn - 0.14

Pb2+

+ 2e- Pb - 0.13

Fe3+

+ 3e- Fe - 0.04

Direction of electron (e-) flow

Anode Zn ZnSO4 (aq) || CuSO4 (aq) Cu Cathode (-)

Direction of electric current flow

(+)


Engineering Chemistry 4

H++ e

- ½ H 0.00

Cu2+

+ 2e- Cu + 0.34

Ag+ + e

- Ag + 0.80

Pt4+

+ 4e- Pt + 0.86

Au+ + e

- Au + 1.69

½ F2 + e- F

- + 2.87

The electro chemical series provide valuable information regarding:

2.1 Relative ease of oxidation or reduction

A system with high reduction potential has a great tendency to undergo reduction. For example, the

standard reduction potentials of F2/F- System and Li

+/Li System is + 2.87V and -3.05V respectively.

The former one can easily gain electrons than the later one. So F2 can easily be reduced to F- and Li is

easily oxidized to Li+.

2.2 Replacement Tendency

Metal with greater oxidation potential can displace metals with lower oxidation potentials from their

salt solution. For Example, Cu2+

has more tendency to replace Zn. Zinc will displace copper from the

solution of CuSO4.

2.3 Predicting spontaneity of redox reactions

Positive value of E of a cell reaction indicates that the reaction is spontaneous. If the value of E is

negative, the reaction is not feasible.

2.4 Calculation of Equilibrium Constant

The standard electrode potential

eqeq

0 KlognF

2.303RTKln

nF

RT E

Hence, 0.0592V

nE

2.303RT

E x nF K log

00

eq

3. ELECTROMOTIVE FORCE (EMF)

Electromotive force is the difference of potential produced by sources of electrical energy, which

can be used to drive current through electrical circuit. It is not a force but a scalar quantity expressed

in volts.

3.1 Electrode Potential

A metal (M) consists of metal ions (Mn+

), with the valence electrons that bind the metal atoms

together. If a metal is in contact with a solution of its own salt, the following two chemical reactions

will take place.



a) Positive metallic ions passing into solution

M Mn+

+ ne- (Oxidation)

b) Positive ions get deposited on the metal electrode

Mn+

+ ne- M (Reduction)

In the first case, „n‟ electrons are left behind on the metal and it acquires negative charge. In

the second case, the metal acquires a positive charge. A dynamic equilibrium is ultimately

established in the reaction. At equilibrium, the potential difference between the metal and solution

attains a constant value, which is called the “Electrode potential”.

Thus, “Electrode potential of a metal is the measure of tendency of a metal to lose or gain electrons,

when it is in contact with a solution of its own salt of unit molar concentration at 25 oC”.

As a convention, the tendency of an electrode to lose electrons is termed as oxidation

potential (+X) and the tendency of an electrode to gain electrons as reduction potential (-X).

3.2 Redox Reaction

Oxidation involves loss of electrons and reduction involves gain of electrons those electrons.

In other words, oxidation and reduction must always go side by side and are not independent in an

electro chemical cell.

If we place zinc metal in a solution of copper sulphate, immediate deposition of Cu takes

place and metal zinc goes into solution as Zn+2

Zn(s) + Cu2+

(aq) Zn2+

(aq) + Cu(s)

In this change, the zinc atom (Zn) is oxidized to Zinc ion (Zn2+

), losing electrons; while the

copper ion (Cu2+

is reduced to copper atom, gaining those electrons.

)Oxidation(2e(aq) Zn Zn(S) 2

Cu(S)(aq) Zn (aq)Cu Zn(S)

Cu(S)2e(aq)Cu22

2 (Reduction)

(Redox)



The overall reaction is called redox or oxidation –reduction reaction. Each of these reactions

is known as a half-reaction. The reaction, in which loss of electrons takes place, is called oxidation-

half reaction; while the other reaction, in which gain of electrons takes place, is called reduction half

reaction.

Thus each electro chemical cell is made up of two electrodes which are dipped in their salt solutions

and connected through a salt bridge and give emf on connection in an external circuit. Each of these

electrodes will have its potential. An electrode dipped in its salt solution is a half cell and its potential

called single electrode potential, which cannot be measured directly.

Nernst equation derivation

3.3 EMF of an Electrochemical Cell

An electrochemical or galvanic cell is obtained by coupling two half cells. Mathematically,

the emf of an electrochemical cell is the algebraic sum of the single electrode potentials; provided

proper signs are given according to the actual reaction taking place on the electrodes.

Standard EMF of the cell = [standard oxidation potential at oxidation half cell reaction + Standard

reduction potential of reduction half cell reaction]

= [standard reduction potential at reduction half cell reaction - Standard

reduction potential of oxidation half cell reaction]

[oxidation potential = - reduction potential]

Ecell = Ecathod – Eanode

Ecell = Eright – Eleft ;

Where Ecell = e.m.f. of the cell

Eright = reduction potential of right hand side electrode

Eleft = reduction potential of left hand side electrode

The positive value of Ecell indicates that the cell reaction is feasible.

Problems on EMF

1. Calculate the e.m.f the following reaction at 25 oC

Eo Zn (Ox) = 0.736 V

Eo Cu (Ox) = -0.337 V



E (Cell) = E(R) - E(L)

= 0.337-(-0.763)

= 0.337 + 0.763

= 1.1volt

Zn Zn2+

aq (1M) Cu2+

aq (0.1M) CuThe potential of the cell at 25 oC is 1. volt. calulate the

potentilal of the cell Zn Zn2+

aq (0.5M) Cu2+

aq (0.01M) Cu at the same temperature

(b)

3.4 Measurement of Electrode Potential

It is not possible to know the

absolute value of a single electrode

potential. We can only measure the

difference in potential between two

electrodes potentiometrically, if we fix

arbitrarily the potential of any one of

electrode. For this purpose, the

potential of a standard hydrogen

electrode (SHE) or normal hydrogen

electrode (NHE) has been arbitrarily

fixed as zero.

The SHE or NHE contains a platinum foil electrode coated with platinum black (for better

adsorption of hydrogen gas) in contact with 1M HCl solution through which hydrogen gas is bubbled

at a constant rate of one atmosphere pressure.

This electrode is represented as..

UNIT II ELECTROCHEMISTRY 2019-20


Pt(s), H2 (g) (1 atm); H+(1 M)

The following reactions take place at this electrode

H+(aq) + e

- → ½ H2 (g) Reduction

½ H2 (g) → H+ (aq) + e

- Oxidation

Thus when both of these reactions can take place at the electrode, it is capable of acting as anode or

cathode depending on the other electrode coupled, to form the cell. The EMF of the cell is measured

with a potentiometer and since the single electrode potential value of the hydrogen electrode is

arbitrarily taken as zero, the single electrode potential of the other electrode can be calculated, since

the EMF value determined is the algebraic sum of the values of single electrode potentials.

In order to measure the electrode potential (for eg; Zn dipped in ZnSO4), this electrode is coupled

with SHE through a salt bridge. Zinc electrode acts as anode, so the electrode potential is oxidation

potential with a value + 0.76V and consequently, its reduction potential is -0.76V. The electrode

potential of copper electrode is +0.34V.

From the above equation, it is clear that

i) If concentration of solution [Mn+

] is increased, the electrode potential increases and vice versa.

ii) If temperature is increased, electrode potential is increased and vice-versa.

When the elements are arranged in increasing order of reduction potential, a series called

electro chemical series is obtained.

4. Batteries

A device that stores chemical energy for a release at a desired later stage as electricity is

known as cell and a series of such cells in a group is called a battery. They are commercial portable

source of electro-chemical cells which supply DC at a constant voltage. They convert chemical energy

of a selected reaction into electrical energy. Batteries can be any of the three types.

4.1. Classification

Primary Battery (or Primary cells): These are the cells in which the cell reaction is not reversible.

When the reactants have for the most part been converted into products, no more electric current is

produced and the battery becomes „dead‟. They can be used as long as the reactants are active and

cannot be recharged. Eg: Leclanche cell

Secondary Battery (or Secondary Cells): Here the cell reaction can be reversed by passing the

direct electric current in opposite direction. A secondary battery may be used through a large number

of cycles of discharging and charging. Eg; Lead storage cell

Flow battery/ Fuel cells: These are the cells in which materials (reactants, products, electrolytes)

pass through the battery, which is simply an electrochemical cell that converts chemical to electrical

energy. Energy can be withdrawn indefinitely as long as fuel supply is maintained. They do not store

energy and cannot be reused. Eg; Hydrogen –oxygen fuel cell



4.2 Dry or Leclanche Cell

This is a cell is without fluid component and hence is conveniently portable. The anode of the cell is

a zinc can, containing an electrolyte consisting of solid NH4Cl, ZnCl2 and MnO2, to which starch is

added to make it thick paste to prevent leakage on transport. A graphite rod acts as a cathode and is

immersed in the electrolyte in the centre of the cell as displayed in the figure below:

The cathode reaction is quite complex.

At Anode: Zn(s) Zn2+

(aq) + 2e-

At Cathode: 2MnO2 + 2NH4Cl (aq) + Zn2+

(aq) + 2e- [Zn(NH3)2]Cl2 (s) + Mn2O3 + H2O

Net reaction: Zn(s) + 2NH4Cl (aq) + 2MnO2 (s) Mn2O3(s) + [Zn(NH3)2]Cl2 (s) + H2O

Since various reactions involved cannot be reversed by passing electricity back through the

cell, this cell is irreversible and for one time use only. It gives a voltage of about 1.5V, which can be

increased further to desired level by connecting such cells in series. Dry cell finds wide variety of

applications in flash-lights, transistor radios, calculators.

Disadvantages

1. When current is drawn rapidly from it, products build up on the electrodes, thereby causing a

drop in the voltage.

2. Since the electrolytic medium is acidic, zinc metal dissolves slowly, thereby cell rundown is

slow.

3. Since the reactants are in acidic medium, zinc undergoes corrosion at a faster rate and decreases

the life of the cell.

4.3 Nickel – Cadmium Battery:

This consists of a cadmium anode and cathode composed of a paste of NiO(OH) (s).

The cell reaction is



Anode: Cd(s) + 2OH-(aq) Cd(OH)2 + 2e

-

Cathode: 2NiO(OH)(s) + 2H2O + 2e- 2Ni(OH)2(s) + 2OH

-(aq)

Net reaction: 2NiO(OH)(s) + Cd(s) + 2H2O(l) Cd(OH)2(s) +2Ni(OH)2(s)

The reaction can be readily reversed, because the reaction products, Ni(OH)2(S) and Cd(OH)2, adhere

to the electrode surface.

It is portable, rechargeable cell with a fairly long life >20 years with a high amperage and maintains a

constant voltage of 1.4 to 1.45 volts.

It can be left for long periods of time without any appreciable deterioration, since no gases are

produced during discharging or charging.

Uses: They are used in electronic calculators, electronic flash units, cordless electronic shavers,

transistors and other battery powered small tools, emergency power supply and military tools.

4.4 Lead acid battery (lead storage cell)

A storage cell is that which can work both as voltaic cell as well as electrical cell. During its

operation it works as a voltaic cell i.e., it supplies electrical energy due to a chemical energy and it

becomes discharged. Then during its charging, it works as an electrolytic cell i.e., chemical is

produced by supplying g of electrical energy. It is also called as lead storage battery.

Anode : Lead (Pb)

Cathode : Lead dioxide (PbO2)

Electrolyte : Dil. H2SO4

A lead storage battery consists of 6 cells each producing 2V output. To increase the current

output of each cell, the cathode plates are joined together and the anode plates are also joined together

and connected in series, we get an output of 12V as shown in figure. ….

Electrode reactions during discharge

At anode : Pb → Pb2+

+ 2e-

Pb2+

+ SO42-

→ PbSO4

At Cathode : PbO2 + 4H+ + 2e

- → Pb

2+ + 2H2O

Pb2+

+ SO42-

→ PbSO4 .

Net reaction: Pb + PbO2 + 4H

+ + 2SO4

2- → 2PbSO4 + 2H2O

Electrode reactions during charging

At anode : PbSO4 + 2e- → Pb + SO4

2-

At Cathode : PbSO4 + 2H2O → PbO2 + SO42-

+ 4H+ + 2e

-

Net reaction: 2PbSO4 + 2H2O → Pb + PbO2 + 4H

+ + 2SO4

2-

The PbSO4 forms during discharge is a solid and sticks to the electrode s. So, it is in a

position to gain or receive the electrons during electrolysis.

Such type of cell is generally used for electrical vehicles, Automobiles, railways, hospitals,

power stations, in telephone exchange, UPS system etc.

4.6 Lithium ion cells (Lithium ion batteries)

These are the batteries in which Lithium act as a cathode and porous carbon as an anode.



During discharging the current flows within the battery from the anode to cathode; Li ions moves

form anode to cathode through the non- aqueous electrolyte.

During the charging, an external power source, the current to pass in the reverses direction. The

positive terminal of the battery connected to the cathode of the battery and negative terminal

is connected to the anode.

Anode : Hard carbon (LiC6), Graphite

Cathode : LiCoO2, LiMnO4, LiNiO2

Electrolyte : Liquid (non- aqueous) electrolyte which contains lithium salts such as LiPF6,

LiBF4 in an organic solvent, such as ethylene carbonate.

The anode half-cell reaction:

LixC6 xLi+ + xe

- + 6C

The cathode half-cell reaction:

xLi+ + xe

- + Lin-xCoO2 LiCoO2

In Lithium- ion battery the lithium ions are transported to and from the cathode or anode, with the

transition metal, Co in Li1-x CoO2 being oxidized from Co3+

to Co4+

during charging , and reduce from

Co4+

to Co3+

during discharge.

Applications:

Used in cell phones, laptops, electric equipment‟s etc.

5. Fuel Cells

A fuel cell is an electro-chemical cell which converts the chemical energy of an easily available

fuel -oxidant system directly into electricity. The chemical energy is provided by the fuel and the

oxidant is stored outside the cell and maintains a continuous supply. The essential process in a fuel

cell is

Fuel + oxygen Oxidation products + electrical energy

Characteristics of a fuel cell

They do not store chemical energy. Reactants are to be supplied constantly, while products are

removed constantly. A fuel cell resembles an engine than a battery. The efficiency of the fuel

cell is about twice that of a conventional power plant for generating electricity. Fuel cell generators

are free from noise, vibration, heat transfer thermal pollution and other problems normally associated

with conventional power plants.

Despite these characteristics, the fuel cells are not yet commercialized in operation. A major problem

lies in the choice and availability of suitable catalysts (for electrodes) able to function efficiently for

long periods to time without deterioration and contamination.

5.1 H2-O2 Fuel cell : One of the simplest and most successful fuel cell is hydrogen–oxygen fuel cell,

which consists essentially of an electrolytic solution such as 25% KOH solution and two inert porous



electrodes. Hydrogen and oxygen gases are bubbled, through the anode and the cathode

compartment.

The following reactions take place.

Anode: 2H2 (g) + 4OH-

(aq) 4H2O + 4e-

Cathode: O2 (g) + 2H2O + 4e- 4OH

- (aq)

Net reaction: 2H2 (g) + O2 (g) 2H2O

The standard emf of the cell is

E0 = E

0ox+E

0Red = 0.83 + 0.40 = 1.23V.

In actual practice, the emf of the cell is 0.8 to 1.0V.

The product discharged by the cell is water.

Usually, a large number of these cells are stacked

together in series to make a battery, called the fuel

cell battery or fuel battery.

The electrodes must meet the stringent requirements. They must

i) Be good conductors

ii) Be good electron sources or sinks

iii) Not to be consumed or deteriorated by the electrolyte heat or electrode reactions.

iv) Be excellent catalysts for the reactions that take place on their surface.

When hydrogen is used as the fuel, the electrodes are made of either graphite impregnated with finely

divided platinum, or a 75:25 alloy of Palladium and Silver or nickel. Electrolytes used are aqueous

KOH or H2SO4 or ion-exchange resin saturated with water. For low temperature operating fuel battery

(-54 oC to 72

oC), potassium thiocyanate dissolved in liquid ammonia is employed.

Applications: Hydrogen – Oxygen fuel cells are used as auxiliary energy source in space vehicles,

submarines and other military vehicles. It has light weight and the product out come is useful water.

The initial cost is high but the maintenance cost is low.

5.2 i) Methyl alcohol – oxygen fuel Cell

The half reactions are

Anode: CH3OH + H2O CO2 (g) + 6H+

(aq) + 6e-

Cathode: 23 O2 (g) + 6H

+ + 6e

- 3H2O

Net reaction: CH3OH + 23 O2 (g) CO2 (g) + 2H2O

ELECTROCHEMISTRY AND APPLICATIONS SYLLABUS Part A · The emf of the cell is 1.1 volts. Direction of...

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