1587884932985_final book 26-04-20

Short Answer Questions

1. Explain principle behind emission spectroscopy.
2. What is meant by the term “Absorption spectroscopy”?
3. What are the components used in Flame Photometry?
4. Give applications of Flame Photometer.
5. Give the application of Fluorescence in Medicines.
6. What are Singlet and Triplet states?
7. What one out of benzene and anthracene will have higher wavelength in

spectroscopy?
8. Which one out of 1,3-pentadiene and 1,4-pentadiene will have higher wavelength in

spectroscopy?
9. Give the table for colour emitted by different elements in Flame Photometry.
10. What are limitations of microwave spectra?

Long Answer Questions

1. Explain Principle, Theory, and Working of Flame Photometer.
2. Distinguish between Fluorescence and Phosphorescence.
3. Give applications of Fluorescence in various fields.
4. Explain Radiative and Non-Radiative Transitions, Intersystem Crossing, Internal

Conversion, Vibrational Relaxation using appropriate Jablonski Diagram.
5. Give advantages and disadvantages of Flame Photometry.
6. In calorimetric estimation of each meta ion in solution, a particular filter is selected.

Give reason.
7. Complexing agents are added to colourless solution of given species during

absorption spectroscopy. Explain.
8. Give a schematic diagram of Flame Photometer. Give the various reactions taking

place.
9. Explain Radiative and Non-Radiative Transitions, Singlet and Triplet states in detail,

using Jablonski Diagram.
10. Discuss DCP, ICP, Spark and Arc Emission Sources.

REFERENCES

A.J. Baker and T. Cairns, Spectroscopic Techniques in Organic Chemistry, Heyden
London, 1965.Google Scholar

E.F.H. Britlain, W.O. George and C.H.J. Wells, Introduction to Molecular
Spectroscopy, Academic Press, London, 1970.Google Scholar

W.G. Richards and P.R. Scott, Structure and Spectra of Atoms, Wiley Eastern Ltd.,
1978.Google Scholar

link.springer.com

https://www.sciencedirect.com/topics/biochemistry-genetics-and- molecular
biology/emission-spectroscopy
http://elchem.kaist.ac.kr/vt/chem-ed/spec/atomic/aes.htm
http://vlab.amrita.edu/?sub=2&brch=193&sim=1351&cnt=1
https://www.horiba.com/en_en/technology/measurement-and- control-
techniques/spectroscopy/what-is-fluorescence-spectroscopy/
https://www.tandfonline.com/doi/abs/10.1080/05704920903435599?journalCo
=laps20
https://wiki.kidzsearch.com/wiki/Spectrometer
https://www.edinst.com/blog/jablonski-diagram/
https://www.sciencedirect.com/science/article/pii/B0122266803001605
https://www.sciencedirect.com/topics/chemistry/radiative-transition
https://www.sciencedirect.com/topics/chemistry/nonradiative-transition
https://scholar.google.co.in/scholar?q=radiative+and+non+radiative+transitions
+in+spectroscopy&hl=en&as_sdt=0&as_vis=1&oi=scholart
https://link.springer.com/chapter/10.1007%2F978-1-4757-6266-2_3
https://www.intechopen.com/books/photon-counting-fundamentals-and-
applications/application-of-fluorescence-spectroscopy-for-microbial-detection-
to-enhance-clinical-investigations
https://www.ncbi.nlm.nih.gov/pmc/articles/PMC2787503/
https://www.researchgate.net/topic/Flame-Photometry
https://www.chromedia.org/chromedia?waxtrapp=mkqjtbEsHiemBpdmBlIEcCAr
B&subNav=cczbdbEsHiemBpdmBlIEcCArBP
https://link.springer.com/chapter/10.1007/978-1-4684-1521-6_7
https://www.studyandscore.com/studymaterial-detail/flame-photometer-
principle-components-working-procedure-applications-advantages-and-
disadvantages
https://www.sanfoundry.com/engineering-chemistry-questions-answers-
molecular-spectroscopy/
https://www.khanacademy.org/test-prep/mcat/physical-processes/light-and-
electromagnetic-radiation-questions/e/light-and-electromagnetic-radiation-
questions
https://www.khanacademy.org/test-prep/mcat/physical-processes/light-and-
electromagnetic-radiation-questions/v/electromagnetic-waves-and-the-
electromagnetic-spectrum
https://www.livescience.com/38169-electromagnetism.html
https://www.sanfoundry.com/engineering-chemistry-questions-answers-

molecular-spectroscopy/
https://www.sanfoundry.com/engineering-chemistry-questions-answers-
electronic-spectroscopy/
TEXTBOOKS
Fundamentals of Molecular Spectroscopy by C.N. Banwell, 4th Edition
Engineering Chemistry by Jain and Jain, 18th edition
Spectroscopy of Organic Compounds by P.S. Kalsi

3.ELECTROCHEMISTRY

3.1) Introduction:

3.1.1) Definition

Electrochemistry is defined as that branch of chemistry which deals with the study of the
production of electricity from the energy that is released during spontaneous chemical
reactions and the use of electrical energy to bring about non-spontaneous chemical
transformations.

There are various places where electrochemistry is important, e.g. electrochemistry is used for
the purpose of electroplating, electrochemistry is required in the production of metals like Na,
Mg. Ca and Al., in the purification of metals and also in Batteries and cells used in various
instruments.

3.1.2) History (From Volta to Faraday):

The story of electrochemistry begins with Alessandro Volta, who announced his invention of the
voltaic pile, the first modern electrical battery, in 1800. The voltaic pile produced a continuous
current and thus opened two new areas of study: the chemical production of electricity and the
effects of electricity on chemicals.

(Alessandro Volta) (Michael Faraday)

Sir Humphry Davy of the Royal Institution in London was one of the most important
experimenters with the new voltaic battery, He realized that the production of electricity by the
voltaic pile depended on the occurrence of chemical reactions, not just on the contact of
different kinds of metals, as Volta had thought.

Davy’s student and successor, Michael Faraday, pursued the relationship between electricity
and magnetism. His two laws of electrochemistry, published in 1834, predict how much product
results from passing a certain amount of current though a chemical compound or its solution, a
process that he named “electrolysis.”

3.2) Electrode Potential

When a metal is placed in a solution of its ions, the metal acquires either a positive or negative
charge with respect to the solution. On account of this, a definite potential difference is
developed between the metal and the solution. This potential difference is called electrode
potential.

For example, when a plate of zinc is placed in a solution having Zn2+ ions, it becomes negatively
charged with respect to solution and thus a potential difference is set up between the zinc plate
and the solution. This potential difference is termed the electrode potential of zinc.

(a) Oxidation: Metal ions pass from the electrode into solution leaving an excess

of electrons and thus a negative charge on the electrode.

The conversion of metal atoms into metal ions by the attractive force of polar water molecules.
M → Mn + ne-

The metal ions go into the solution and the electrons remain on the metal making it negatively
charged. The tendency of the metal to change into ions is known as electrolytic solution pressure.

(b) Reduction: Metal ions in solution gain electrons from the electrode leaving a

positive charge on the electrode. Metal ions start depositing on the metal surface leading to a

positive charge on the metal. Mn+ + ne- → M

In the beginning, both these changes occur with different speeds but soon an equilibrium

is established. M Mn+ + ne-

In practice, one effect is greater than the other,

If first effect is greater than the second, the metal acquires a negative charge with respect to

solution and

If the second is greater than the first, it acquires positive charge with respect to solution, thus in

both the cases a potential difference is set up.

The magnitude of the electrode potential of a metal is a measure of its relative tendency to lose

or gain electrons, i.e., it is a measure of the relative tendency to undergo oxidation (loss of

electrons) or reduction (gain of electrons).

3.2.1) Types of electrode potential:

Depending on the nature of the metal electrode to lose or gain
electrons, the electrode potential may be of two types:

• Oxidation potential: When electrode is negatively charged with respect to solution, i.e., it

acts as anode. Oxidation occurs. M → Mn+ + ne-

• Reduction potential: When electrode is positively charged with respect to solution, i.e., it

acts as cathode. Reduction occurs.
Mn+ + ne- → M

It is not possible to measure the absolute value of the single electrode potential directly. Only
the difference in potential between two electrodes can be measured experimentally. It is,
therefore, necessary to couple the electrode with another electrode whose potential is known.

This electrode is termed as a reference electrode. The EMF of the resulting cell is measured
experimentally. The EMF of the cell is equal to the sum of potentials on the two electrodes.
Emf of the cell = EAnode + ECathode = Oxidation potential of anode + Reduction potential of
cathode
Knowing the value of reference electrode, the value of other electrode can be determined.

3.2.2) Classification of electrodes

1. Metal/metal ion electrode:

If a metal plate is immersed into a solution of same metal salt ,a potential is
developed surrounding the metal plate.The electrode of this type is called 1st kind electrode
electrode,eg

Zn|Zn^+2(aq) The
electrode reaction is :

Zn→Zn^+2 +2e
The electrode potential is given by: ɸ

Zn=ɸ0 Zn-RT/nF ln aZn+2 [aZn=1]

2. Metal/sparingly soluble/salt saturated salt solution:

A metal plate,coated with its sparingly soluble salt,is immersed in a solution
containing the same sparingly soluble salt(soluble) and an easily soluble salt containing the
same anion.One example is:

The electrode reaction is:
Ag→Ag+ +e equation 1
Ag+ +Cl-→AgCl(s) equation 2
Adding equation 1 and 2;
Ag +Cl→AgCl(s) +e The electrode potential is given

by:
ɸAg|AgCl(s)=ɸ0Ag|AgCl(s) - RT/nF ln a AgCl/a Ag(s). aCl
ɸAg|AgCl(s)=ɸ0Ag|AgCl(s) + RT/nF ln a AgCl/a Ag(s). aCl

3. Redox electrode:
If a platinum is immersed into a solution containing a multivalent metal ion,redox

electrode is formed.Pt-plate is chosen since it does not participate in chemical reaction but
serves only as carrier of electrons.One example of redox electrode is Ptplate dipped into a
solution containing both Fe+2 and Fe+3 ions.The electrode reaction is: Fe+2→ Fe+3 +e

The electrode potential is given by:
ɸFe+2|Fe+3=ɸ0Fe+2|Fe+3 - RT/F ln aFe+3/aFe+2

Let’s test your knowledge

Q1) The total potential is the sum of the potentials produced by the reactions_____

a) in the solution
b) at the salt bridge
c) at the two electrodes
d) at the anode only

Q2) Another name for electrical potential is

a) electrical force
b) electromotive force
c) electrolytic force
d) electrochemical force

Q3) The electromotive force (EMF) of a cell equals

a) EMF oxidation − EMF reduction
b) EMF reduction − EMF oxidation
c) EMF oxidation + EMF reduction
d) EMF reduction + EMF oxidation

Q4) Consider a voltaic cell based on the redox reaction:

With the following reduction potentials:

Which species is oxidized?

a) Hg 2+
b) Hg
c) Cu
d) Cu 2+

Q5) With the help of the data of the above question. What is the EMF of the cell?

a) 0.34 V
b) 0.92 V
c) 0.58 V
d) -0.58V

Anwers

Q1 Q2 Q3 Q4 Q5

cbac c

3.3 Concept of Standard Electrode Potential:

The potential of an electrode cannot be accurately measured. A reference electrode has a
known electrode potential and is stable. Its high stability is achieved by employing the redox
system, which must contain saturated concentrations in each of the participating solutions of
the reaction.

3.3.1 Reference Electrode

Uses of reference electrodes are numerous, but the most important of all is in the
electrochemical cell. This is where it's used as a half cell in the electrochemical cell to allow for
the determination of the other half's cell potential. It is not feasible to connect measuring
device to the solution of the electrode as it may itself create an equilibrium along with the
existing one which in turn will give wrong value. This problem can be overcome if we use
reference electrode.

3.3.2 Types of Reference Electrode

Reference electrodes can be classified as aqueous, calomel, non-aqueous and own constructing.
For an electrode to be termed as reference electrode it must follow the follow criteria:

1. The potential of the electrode must be known, the conditions of utility.

2. The potential of the electrode must show minimum variation with temperature or other
physical factors.

The most common aqueous reference electrodes used include:
· Standard hydrogen electrode
· Normal hydrogen electrode
· Saturated calomel electrode
· Reversible hydrogen electrode
· Silver chloride electrode
· Copper-copper sulfate electrode
· PH electrode
· Dynamic hydrogen electrode
· Palladium-hydrogen electrode

3.3.3) Standard Hydrogen Electrode (SHE):

• What is it?

The Standard Hydrogen Electrode is often abbreviated to SHE, and its standard electrode
potential is declared to be 0 at a temperature of 298K. This is because it acts as a reference for
comparison with any other electrode .

• Half Cell Reactions:

The redox half cell of the SHE is where the following reaction takes place:
2H+ (aq) + 2e– → H2 (g)
The reaction given above generally takes place on a platinum electrode. The pressure of the
hydrogen gas present in this half cell equals 1 bar.

• Why platinum?

Platinum is used in the Standard Hydrogen Electrode due to the following reasons:
• Platinum is a relatively inert metal which does not corrode easily.
• Platinum has catalytic qualities which promotes the proton reduction reaction.
• The surface of platinum can be covered with platinum black, a fine powder of platinum.

This type of platinum electrode is called a platinized platinum electrode.

• Construction:

A platinum electrode which is covered in finely powdered platinum black (platinized platinum
electrode).A solution of acid having a H+ molarity of 1 mole per cubic decimeter.The SHE also
contains a hydroseal which is used to prevent the interference of oxygen.The other half-cell of
the entire Galvanic cell must be attached to the Standard Hydrogen Electrode through a
reservoir in order to create an ionically conductive path. This can be done through a direct
connection, through a narrow tube, or even through the use of a salt bridge.
The platinized platinum surface has a very high adsorption activity. Therefore, this surface must
be protected from atmospheric oxygen as well as from organic substances. Substances such as
arsenic and sulfur compounds can deactivate or poison the catalyst.

Advantages of Standard Hydrogen Electrode (S.H.E.):

• Small potential is developed on the hydrogen electrode, hence it can be taken as zero.
• In determining the single electrode potential, using S.H.E. as a reference, the potential

of the unknown potential will be equal to the e.m.f. of the cell.

Disadvantages of Standard Hydrogen Electrode (S.H.E.):

• It is not convenient to assemble the apparatus.
• It is difficult to maintain the pressure of hydrogen gas and concentration of HCl.
• It is difficult to get pure, dry hydrogen gas and prepare ideal platinised platinum plate.
• The impurities present in H2 and HCl poison the Pt, and affect the equilibrium at the

electrode.

Example:

Determination of Standard Electrode Potential of Zn/Zn2+ Electrode
A zinc rod is dipped in 1 M zinc sulphate solution. This half-cell is combined with a standard
hydrogen electrode through a salt bridge. Both the electrodes are connected with a
voltmeter. The deflection of the voltmeter indicates that current is flowing from hydrogen
electrode to metal electrode or the electrons are moving from zinc rod to hydrogen electrode.
The zinc electrode acts as an anode and the hydrogen electrode as cathode and the cell can
be represented as

Cell Representation

Oxidation half Reduction half
reaction reaction

|| 2H(aq)|

Zn|Zn2+ (aq)/Anode(-) Zn → Zn2+ + 2H+ + 2e-
H2 (g)/Cathode (+) 2e- → H2↑

The EMF of the cell is 0.76 volt
ECell = EoAnode + EoCathode
0.76 = EoAnode + 0 or EoAnode = +0.76 V
As the reaction on the anode is oxidation, i.e.,
Zn → Zn2+ + 2e-
EoAnode is the standard oxidation potential of zinc. This potential is given the positive sign.

Eoox (Zn/Zn2+) = +0.76 volt
So standard reduction potential of Zn, i.e., Eo (Zn/Zn2+)
= Eoox = -(+0.76)
= -0.76 volt
The EMF of such a cell gives the positive value of standard oxidation potential of metal M. The
standard reduction potential (Eo) is obtained by reversing the sign of standard oxidation
potential.

The standard hydrogen half-cell paired with a zinc half-cell.

3.3.4) Calomel Electrode:

What is it?

The saturated calomel electrode is a reference electrode based on the reaction between
elemental mercury and mercury(I) chloride. Calomel electrode is the mercurymercurous
chloride electrode. It has been widely replaced by the silver chloride electrode, however the
calomel electrode has a reputation of being more robust.

Half Cell Reaction:

The calomel electrode can act as anode or cathode depending on the nature of other electrode
of the cell.
When it acts anode, the electrode reaction is

2Hg(l) → Hg22+ + 2e-
Hg22+ + 2Cl- → Hg2Cl2
--------------------------------------------------------------- 2 Hg + 2Cl- → Hg2Cl2
+ 2e- (Oxidation reaction)

When it acts as cathode, the electrode reaction is,
Hg22+ + 2e- → 2 Hg
Hg2Cl2 → Hg22+ + 2Cl-

________________________________________
Hg2Cl2 + 2e- → 2Hg (l) + 2 Cl- [Reduction reaction]

Construction:

It consists of glass vessel having bent side tube. Pure mercury is placed at the bottom of the
tube. Which is covered with a paste of mercury- mercurous chloride (Hg+Hg2Cl2) i.e., calomel.
The remaining portion of the cell is filled with a solution of normal (1N) or decinormal (0.1N) or
saturated KCl. A platinum wire sealed into a glass tube is dipped into mercury layer is used to
provide the external electrical contact. The side tube is used for making electrical contact with a
salt bridge.

The electrode can be represented as Hg (l)

| Hg2Cl2 (s) | Cl-

The net reversible electrode reaction is,

Hg2Cl2(s) + 2e- --------> 2 Hg (l) + 2 Cl-

Nerst Equation:

Electrode potential is given by
E = Eo – 2.303 RT/2F log [Cl-]2

= Eo - 0.0591 log [Cl-] at 298 K
The electrode potential is decided by the concentration of chloride ions & the electrode is reversible with
chloride ions at 298K, the electrode potentials are as follows.
0.1N KCl electrode (0.334V)
1 N KCl electrode (0.281 V)
Saturated KCl electrode (0.2422 V)

Advantages of Standard Calomel Electrode:

• It is very handy, compact and easy to transport.
• Its potential can remain constant and it can easily be reproduced.
• It is easy to construct and maintain.

Let’s Test your Knowledge:

Question 1: Which of the following equations represents
calomel electrode?
a. Hg2Cl2 + e- Hg + Cl-
b. AgCl (s) + e– → Ag + Cl–
c. Fe3+ + e– → Fe2+
d. H2 2H+ + 2e-
Question 2: Magnitude of electrode potential does not depend on.
a. Pressure
b. Temperature
c. Concentration
d. Nature of electrode

Question 3: In which of the following elements is not a component of SHE?
a. Platinum
b. Hydrogen
c. Mercury
d. Gold
Question 4: SHE can be used as..
a. only anode
b. only cathode
d. both anode and cathode
c. neithet anode nor cathode
Question 5: E° value for Cu2+/Cu volt is a. 0.34 V
b. 0.58 V
c. 1.02 V

d. 1.51 V

ANWERS:

Q.1 Q.2 Q.3 Q.4 Q.5

aa dda

3.4) Electrochemical cell:

An electrochemical cell is a device capable of either generating electrical energy from
chemical reactions or using electrical energy to cause chemical reactions. There are diff types
of electrochemical cells: i. Galvanic Cell ii. Electrolytic iii. Fuel cell iv. Concentration cell

3.5) TYPES OF ELECTROCHEMICAL CELLS:

3.5.1) Voltaic cell (Galvanic cell):

1. In galvanic cell the free energy change of electrode redox reaction is converted into
electric energy; that is del G=-nFE.

2. The free energy change in a galvanic cell must be negative and the potential difference
of the electrode (cell voltage) must be positive.

3. Energy is released from spontaneous redox reaction. System does work on
load/surrounding.

4. The anode of a galvanic cell is negatively charged, since the spontaneous oxidation at
the anode is the source of the cell's electrons or negative charge. The cathode of a
galvanic cell is its positive terminal.

5. In both galvanic and electrolytic cells, oxidation takes place at the anode and electrons
flow from the anode to the cathode.

6. The redox reaction in a galvanic cell is a spontaneous reaction. 7. For this reason,
galvanic cells are commonly used as batteries.

8. Galvanic cell reactions supply energy which is used to perform work.
9. The energy is harnessed by situating the oxidation and reduction reactions in separate

containers, joined by an apparatus that allows electrons to flow.

The examples of galvanic cells are:
• Primary irreversible cells (dry cells)
• Secondary cells (lead storage batteries)

Setup of a Galvanic Cell

• In order to create a galvanic cell, one would have to go through the following setup.
The cell would ideally include two electrodes. One of these electrodes, the cathode,
shall be a positively charged electrode while the other, shall be the anode, the
negatively charged electrode.

• These two electrodes shall form the two essential components of the galvanic cell.
The chemical reaction related to reduction shall take place at the cathode while
the oxidation half-reaction shall take place at the anode. As has already been said,
any two metals can be used to create the chemical reaction.

Where is Galvanic cell used?

• A common galvanic cell is the Daniel cell. Simple cells are galvanic cells.
They are used from TV remotes to car batteries, from clocks to inverters.

3.5.2) Electrolytic cell:

1. It is possible to construct a cell that does work on a chemical system by driving an
electric current through the system. These cells are called electrolytic cells.

2. It is a device in which electrical energy from an external source can be used to
produce chemical reactions.

3. Energy is absorbed to drive non-spontaneous redox reaction. Surroundings (power
supply) do work on system(cell).

4. In electrolytic cells ,the reaction is being driven in the non-spontaneous direction by
external electrical force and free energy change is positive.

5. Electrolytic cells, like galvanic cells, are composed of two half-cells--one is a reduction
half-cell, the other is an oxidation half-cell.

6. The direction of electron flow in electrolytic cells, however, may be reversed from
the direction of spontaneous electron flow in galvanic cells, but the definition of both
cathode and anode remain the same, where reduction takes place at the cathode
and oxidation occurs at the anode. Because the directions of both half-reactions have
been reversed, the sign, but not the magnitude, of the cell potential has been
reversed.

Set up of an electrolytic cell

• Electrolytic cells are very similar to voltaic (galvanic) cells in the sense that both require a
salt bridge, both have a cathode and anode side, and both have a consistent flow of
electrons from the anode to the cathode.

Where is Electrolytic cell used?

• The most common form of Electrolytic cell is the rechargeable battery (cell phones, etc) or
electroplating. While the battery is being used in the device it is a galvanic cell function
(using the redox energy to produce electricity). While the battery is charging it is an
electrolytic cell function (using outside electricity to reverse the completed redox reaction).

• Commercially, electrolytic cells are used in electrorefining and electrowinning of several
non-ferrous metals.

• Almost all high-purity aluminium, copper, zinc and lead is produced industrially in
electrolytic cells.

3.5.3) Concentration cell:

1. A concentration cell is also an electrochemical device that generates electrical energy
when two electrodes of the same metal are in contact with solutions of it‟s ions at
different concentrations.

2. In a concentration cell, too, the free energy change of electrode reactions is converted into
electric energy.

3. As the cell as a whole strives to reach equilibrium, the more concentrated half cell is diluted
and the half cell of lower concentration has its concentration increased via the transfer of
electrons between these two half cells. Therefore, as the cell moves towards chemical
equilibrium, a potential difference is created.

A detailed diagram of a concentration cell and its discharge process is given below.

Types of Concentration Cells

Concentration cells can be classified into two types, namely:
• Electrode Concentration Cells
• Electrolyte Concentration Cells

1.Electrode Concentration Cells

 These cells consist of identical solutions used as electrolytes in each half-cell. However,
the half-cells differ in the concentration of the electrode (the electrodes are made up
of the same material).

 An Example for this type of cell would be a cell consisting of two hydrogen electrodes
which are subjected to varying pressures but are immersed in the same solutions
(containing hydrogen ions).

2.Electrolyte Concentration Cells

 These cells consist of identical electrodes immersed in the solutions of the same
electrolytes but with varying concentrations. In these cells, the electrolyte tends to
diffuse from higher concentration solutions towards solutions of lower concentration.

 An example for this type of cell is a cell where the anode consists of
Zn/Zn2+(0.1M) whereas the cathode consists of Zn2+(0.01M)/Zn. In this cell, the flow of
electrons from the anode to the cathode is due to the reduction of Zn2+ ions at the
cathode into metallic zinc.

Where concentration cell used?

• The concentration cells are used to determine the solubility of sparingly soluble salts,
valency of the cation of the electrolyte and transition point of the two allotropic forms
of a metal used as electrodes, etc.

3.5.4) Fuel cell:

1. A fuel cell is also an electrochemical device, which operates with continuous replenishment
of the fuel at the electrode and no charging is required .In a fuel cell, the free energy
change of electrode redox reactions is converted into electrical energy.

Set up:

 A fuel cell needs three main components to create the chemical reaction:
an anode, cathode and an electrolyte. First, a hydrogen fuel is channeled to the

anode via flow fields. Hydrogen atoms become ionized (stripped of electrons), and now
carry only a positive charge. Then, oxygen enters the fuel cell at the cathode, where it
combines with electrons returning from the electrical circuit and the ionized hydrogen
atoms. Next, after the oxygen atom picks up the electrons, it then travels through the
electrolyte to combine with the hydrogen ion. The combination of oxygen and ionized
hydrogen serve as the basis for the chemical reaction.
 A polymer electrolyte membrane permits the appropriate ions to pass between the
anode and the cathode. If the electrolyte gave free control for all electrons or ions to
pass freely, it would disrupt the chemical reaction. At the end of the process the
positively charged hydrogen atoms react with the oxygen to form both water and heat
while creating an electrical charge.
 Within the fuel market there are many different applications with different power
requirements. In order to provide adequate power, individual fuel cells can be
assembled together forming a stack. A fuel cell stack can be sized for just the right
amount of energy for the application.

Where are fuel cells used?

Fuel Cells are used in both stationary and motive power applications for:
• Cars, trucks, buses, and recreational vehicles
• Material handling equipment
• Act as a primary power source for high-volume data centers or commercial, industrial,
and residential buildings
• Backup power source to critical computer and communications networks Generating
power on-site

Let’s Test Your Knowledge

Q1) Galvanic cells are also named as
A. electrolytic cells
B. battery cells
C. Daniel cells
D. John cells

Q2) Voltaic cells generate electricity by
A. spontaneous redox reaction
B. non spontaneous redox reaction
C. sublimation reaction
D. thermochemical reaction

Q3) In a fuel cell, the positive ions (hydrogen ions) are produced at the _____.
A. Anode
B. Electrolyte Center
C. Cathode
D. Negative battery terminal

Q4) Voltaic cells, such as those diagrammed in your text, the salt bridge _______ .

A. is not necessary in order for the cell to work
B) acts as a mechanism to allow mechanical mixing of the solutions

C) allows charge balance to be maintained in the cell
D) is tightly plugged with firm agar gel through which ions cannot pass
E) drives free electrons from one half-cell to the other

Anwers Q1-C Q2- A Q3- A Q4- C

3.6) ELECTROCHEMICAL SERIES

By measuring the potentials of various electrodes versus standard hydrogen electrode (SHE), a
series of standard electrode potentials has been established.

When the electrodes (metals and non-metals) in contact with their ions are arranged on the
basis of the values of their standard reduction potentials or standard oxidation potentials, the
resulting series is called the electrochemical or electromotive or activity series of the elements.
By international convention, the standard potentials of electrodes are tabulated for reduction
half reactions, indicating the tendencies of the electrodes to behave as cathodes towards SHE.
Electrodes with positive E° values for reduction half reactions do in fact act as cathodes versus
SHE, while those with negative E° values of reduction half reactions behave instead as anodes
versus SHE. The electrochemical series is shown in the following table.

Standard Aqueous Electrode Potentials at 25°C 'The Electrochemical Series‟

Element Electrode Reaction (Reduction)

Ni 0Ni2+ + 2e- → Ni -0.25

Sn Sn2+ + 2e- → Sn -0.14

H2 2H+ + 2e- → H2 0.00

Cu Cu2+ + 2e- → Cu +0.337

I2 I2 + 2e- → 2I- +0.535

Ag Ag+ + e- → Ag +0.799

Hg Hg2+ + 2e- → Hg +0.885
Br2 Br2 + 2e- → 2Br- +1.08

Cl2 Cl2 + 2e- → 2Cl- +1.36
Au3+ + 3e- → Au +1.50

Au

F2 F2 + 2e- → 2F- +2.87

The negative sign of standard reduction potential indicates that an electrode when joined with
SHE acts as anode and oxidation occurs on this electrode.
For example, standard reduction potential of zinc is -0.76 volt.
When zinc electrode is joined with SHE, it acts as anode (-ve electrode) i.e., oxidation occurs on
this electrode. Similarly, the +ve sign of standard reduction potential indicates that the electrode
when joined with SHE acts as cathode and reduction occurs on this electrode.

3.6.1) Characteristics of Electrochemical Series

• The substances which are stronger reducing agents than hydrogen are placed above
hydrogen in the series and have negative values of standard reduction potentials.

• All those substances which have positive values of reduction potentials and placed below
hydrogen in the series are weaker reducing agents than hydrogen.

• The substances which are stronger oxidising agents than H+ion are placed below hydrogen in
the series.

• The metals on the top (having high negative values of standard reduction potentials) have
the tendency to lose electrons readily. These are active metals.

• The activity of metals decreases from top to bottom.
• The non-metals on the bottom (having high positive values of standard reduction potentials)

• have the tendency to accept electrons readily. These are active non-metals.
• The activity of non-metals increases from top to bottom.

3.6.2) Applications of Electrochemical Series

1.) Reactivity of Metals

The activity of the metal depends on its tendency to lose electron or electrons, i.e., tendency to
form cation (M"+). This tendency depends on the magnitude of standard reduction potential. The
metal which has high negative value (or smaller positive value) of standard reduction potential
readily loses the electron or electrons and is converted into cation. Such a metal is said to be
chemically active. The chemical reactivity of metals decreases from top to bottom in the series.
Metals like Fe, Pb, Sn, Ni, Co, etc., which lie a little down in the series do not react with cold water
but react with steam to evolve hydrogen. Metals like Cu, Ag and Au which lie below hydrogen are
less reactive and do not evolve hydrogen from water.

2.) Electropositive Character of Metals

The electropositive character also depends on the tendency to lose electron or electrons.
On the basis of standard reduction potential values, metals are divided into three groups:

• Strongly electropositive metals: Metals having standard reduction potential near about 2.0
volt or more negative like alkali metals, alkaline earth metals are strongly electropositive in
nature.

• Moderately electropositive metals: Metals having values of reduction potentials between
0.0 and about -2.0 volt are moderately electropositive. Al, Zn, Fe, Ni, Co, etc., belong to this
group.

• Weakly electropositive metals: The metals which are below hydrogen and possess positive
values of reduction potentials are weakly electropositive metals. Cu, Hg, Ag, etc., belong to
this group.

3.) Displacement Reactions

To predict whether a given metal will displace another, from its salt solution
Displacement of one non-metal from its salt solution by another
Displacement of hydrogen from water

4.) Reducing Power of Metals

Reducing nature depends on the tendency of losing electron or electrons. More the negative
reduction potential, more is the tendency to lose electron or electrons. Thus, reducing nature
decreases from top to bottom in the electrochemical series. The power of the reducing agent
increases as the standard reduction potential becomes more and more negative. Sodium is a
stronger reducing agent than zinc and zinc is a stronger reducing agent than iron.

5.) Oxidising Nature of Non-metals

Oxidising nature depends on the tendency to accept electron or electrons. More the value of
reduction potential, higher is the tendency to accept electron or electrons. Thus, oxidising nature
increases from top to bottom in the electrochemical series. F2 (Fluorine) is a stronger oxidant
than Cl2, Br2 and I2.

6.) Thermal Stability of Metallic Oxides

The thermal stability of the metal oxide depends on its electropositive nature. As the
electropositivity decreases from top to bottom, the thermal stability of the oxide also decreases
from top to bottom. The metals which come below copper form unstable oxides, i.e., these are
decomposed on heating and vice versa.

7.) Products of Electrolysis

In general, in such competition the ion which is stronger oxidising agent (high value of standard
reduction potential) is discharged first at the cathode.
The increasing order of deposition of few cations is:
K+, Ca2+, Na+, Mg2+, Al3+, Zn2+, Fe2+, H+, Cu2+, Ag+, Au3+

8.) Corrosion of Metals

Corrosion is defined as the deterioration of a substance because of its reaction with its
environment.

3.6.4) Mnemonics for electrochemical series:

Let’s Test your Knowledge:

Q 1: Which of the following metals has most negative value of standard electrode reduction
potential?
a. Na
b. Ca
c. Mg
d. K
Q 2: Which of the following metals is least electropositive one? a. Al
b. Cu
c. Fe
d. Zn
Q3: Which of the following metals is most reactive one ? a. Na
b. Ca
c. Mg
d. K
Q 4: Which of the following metals will not displace Cu from aqueous solution of CuSO4?
a. Ag
b. Zn

c. Ni
d. Fe

Q 5: Which of the following metal ions will discharge first at electrode?
a. K+
b. Ca2+
c. Mg2+
d. Na+

3.7) Nernst Equation:

In electrochemistry, the Nernst equation can be used, in conjunction with other information, to
determine the reduction potential of a half-cell in an electrochemical cell. It can also be used to
determine the total voltage, or electromotive force, for a full electrochemical cell. It is named
after the German physical chemist who first formulated it, Walther Nernst.
The Nernst equation gives a formula that relates the electromotive force of a nonstandard cell
to the concentrations of species in solution:

E = EO – [(2.303*R*T)/n*F]*log[M / M+] In
this equation:

• E is the electromotive force of the non-standard cell
• EO is the electromotive force of the standard cell
• n is the number of moles of electrons transferred in the reaction
• R is universal gas constant (8.314 J K-1 mol-1)
• F is Faraday constant (96500 C mol-1)
• log[M / M+] is the natural log of chemical activity coefficients of the relevant species

3.7.1) Derivation

Consider a metal in contact with its own salt aqueous solution. Reactions of metal losing an
electron to become an ion and the ion gaining electron to return to the atomic state are equally
feasible and are in an equilibrium state. Mn+ + ne– → nM
In the reduction reaction, „n‟ moles of an electron is taken up by the ion against a reduction
potential of Ered.
1. The work done in the movement of electron
Wred = nFEred
Where,

• F is Faraday = 96487 coulomb = electrical charge carried by one mole of electrons
2. Change in the Gibbs free energy is an indication of the spontaneity and it is also equal to the

maximum useful work (other than volume expansion) done in a process.
Combining work done and Gibbs free energy change:
Wred = nFEred = – ∆G or ∆G = – nFEred
3. Change in the free energy at standard conditions of 298K and one molar /one atmospheric

pressure conditions is ∆G°. From the above relation, it can be written that ∆G° = – nFE°red
Where,

• E°red is the reduction potential measured at standard conditions.

4. During the reaction, concentration keeps changing and the potential also will decrease with
the rate of reaction.

To get the maximum work or maximum free energy change, the concentrations have to be
maintained the same. This is possible only by carrying out the reaction under a reversible
equilibrium condition.
For a reversible equilibrium reaction, van Hoff isotherm says:
∆G = ∆G° + RT lnK

Where,

• K is the equilibrium constant T is the temperature in Kelvin scale.
• K = Product/Reactant = [M]n/[M]n+
• R is the Gas constant =8 .314J/K mole

5. Substituting for free energy changes in ant Hoff equation,
– nFEred = – nFE°red + RT ln [M]/[Mn+] = – nFE°red + 2.303 RT log [M]n/[Mn+]
Dividing both sides by – nF,
Ered = E°red – 2.303RTnF1log[M]n[Mn+] or,
EMn+/M=EoMn+/M−2.303RTnF1log[M]n[Mn+]
The activity of the metal is, always considered as equal to unity.
Ered = E°red – or EMn+/M
= EoMn+/M- 2.303RTnFlog1[Mnn+]
This relation connecting reduction potential measurable at conditions other than standard
conditions to the standard electrode potential is the Nernst equation.
For reaction conducted at 298K but at different concentrations, Nernst Equation is;
EMn+/M= EoMn+/M- 2.303×8.314×293n96500log1[Mnn+]
= EoMn+/M- 0.0591nlog1[Mnn+]

3.7.2) Nernst Equation Applications

The Nernst equation can be used to calculate:
• Single electrode reduction or oxidation potential at any conditions
• Standard electrode potentials
• Comparing the relative ability as a reductive or oxidative agent.
• Finding the feasibility of the combination of such single electrodes to produce electric
potential.

• Emf of an electrochemical cell
• Unknown ionic concentrations
• The pH of solutions and solubility of sparingly soluble salts can be measured with the

help of the Nernst equation.

3.7.3) Limitations of Nernst Equation

• The activity of an ion in a very dilute solution is close to infinity and can, therefore, be
expressed in terms of the ion concentration. However, for solutions having very high
concentrations, the ion concentration is not equal to the ion activity. In order to use the
Nernst equation in such cases, experimental measurements must be conducted to
obtain the true activity of the ion.

• Another shortcoming of this equation is that it cannot be used to measure cell potential
when there is a current flowing through the electrode. This is because the flow of
current affects the activity of the ions on the surface of the electrode. Also, additional
factors such as resistive loss and overpotential must be considered when there is a
current flowing through the electrode.

3.8) Relation between Corrosion and Electrochemistry:

Any process of deterioration or destruction and consequent loss of a solid metallic material
through an unwanted chemical or electrochemical attack by its environment at its surface
is called corrosion. Thus corrosion is a reverse process of extraction of metals.

One of the main types of corrosion is wet corrosion also known as electrochemical corrosion.
This type of corrosion is observed when

• A conducting liquid is in contact with a metal (or)
• When two dissimilar metals (or) alloys are either immersed (or) dipped partially in

a solution.
• The corrosion occurs due to the existence of separate anodic and cathodic areas or parts

between which current flows through the conduction soln.
• In the anodic area oxidation reaction takes place so anodic metal is destroyed by

dissolving (or) forming a compound such as an oxide.
Hence corrosion always occurs at anodic areas
.: At Anode
M --------> Mn+ + ne–

.: At cathode
Mn+ + ne– --------> M

• In cathodic area, reduction reaction (gain of e– s) takes place. The metal which is acting
as cathode is in its reduced form only. Therefore it cannot be further reduced. Therefore
cathodic reactions do not affect the cathode.

• So at cathodic part dissolved constituents in the conducting medium accept the
electrons to form some ions like OH-, O2-, etc.

• The metallic ions from anodic part and non-metallic ions from cathodic part
diffuse towards each other through conducting medium and form a corrosion product
somewhere between anode and cathode.

• The e- s which are set free at anodic part flow through the metal and are finally
consumed in the cathodic region.

Thus we may sum up that electrochemical corrosion involves:

i)The formation of anodic and cathodic areas.
ii)Electrical contact between the cathodic and anodic parts to enable the conduction of
electrons.
iii)An electrolyte through which the ions can diffuse or migrate this is usually provided
by moisture.
iv)Corrosion of anode only
v)Formation of corrosion product is somewhere in between cathode and anode

3.9) Recent Development in Electrochemistry:

3.9.1) Paper-based Analytical Devices ( PADs)

Research into electrochemical sensing in paper devices has increased dramatically in the past
few years because electrochemical methods allow sensitive, selective, and quantitative
detection once incorporated into paper devices. The electrochemical techniques coupled to
those (paper-based analytical devices) PADs include cyclic voltammetry, amperometry,
coulometry, and potentiometry. The electrochemical systems employed by those PADs involve
microelectrodes, convective mass transfer, surface modified electrodes, flow injection, signal
amplification, ion-selective electrodes, and ion exchange membranes, respectively. This
minireview discusses general aspects of the aforementioned methods and the resulting
electrochemical PADs.

3.9.2) Biosensors

There are different types and different generations of biosensors. A product of reaction diffuses
to the transducer in the first generation biosensors (based on Clark biosensors). The mediated
biosensors or second generation biosensors use specific mediators between the reaction and
the transducer to improve sensitivity. The second generation biosensors involve two steps: first,
there is a redox reaction between enzyme and substrate that is reoxidized by the mediator, and
eventually the mediator is oxidized by the electrode. No normal product or mediator diffusion is
directly involved in the third generation biosensors, direct biosensors. Based on the type of
transducer, current biosensors are divided into optical, mass, thermal, and electrochemical
sensors. They are used in medical diagnostics, food quality controls, environmental monitoring,
and other applications. These biosensors are also grouped under two broad categories of
sensors: direct and indirect detection systems. Moreover, these systems could be further
grouped into continuous or batch operation. Therefore, amperometric biosensors and their
current applications are focused on more in detail since they are the most commonly used
biosensors in monitoring and diagnosing tests in clinical analysis. Problems related to the
commercialization of medical, environmental, and industrial biosensors as well as their
performance characteristics, their competitiveness in comparison to the conventional analytical
tools, and their costs determine the future development of these biosensors.

3.10) Solved Problems on Nernst Equation:

3.10.1) The standard electrode potential of zinc ions is 0.76V. What
will be the potential of a 2M solution at 300K?

Solution:
The Nernst equation for the given conditions can be written as follows;
EMn+/M = Eo – [(2.303RT)/nF] × log 1/[Mn+]
Here,

• E° = 0.76V
• n=2
• F = 96500 C/mole
• [Mn+] = 2 M
• R =8.314 J/K mole
• T =300 K
Substituting the given values in Nernst equation we get,
EZn2+/Zn = 0.76 – [(2.303×8.314×300)/(2×96500)] × log 1/2 = 0.76 – [0.0298 × (-0.301)]
= 0.76 + 0.009 = 0.769V
Therefore, the potential of a 2M solution at 300K is 0.769V.

3.10.2) From the following standard potentials, arrange the metals in
the order of their increasing reducing power.

• Zn2+(aq) + 2e– → Zn(s): E° = -0.76 V
• Ca2+(aq) + 2e– → Ca(s): E° = -2.87 V
• Mg2+(aq) + 2e– → Mg(s): E° = -2.36 V
• Ni2+(aq) + 2e– → Ni(s): E° = -0.25 V
• Ni(s) → Ni2+(aq) + 2e– : E° = +0.25 V
Reducing power of a metal increases with its ability to give up electrons ie lower standard
potentials. Arranging the reduction potentials in the decreasing order gives the increasing order
of reducing power of metals.
Increasing order of reduction potentials is Ni (-0.25V) < Zn (-0.76V) < Mg(-2.36V) < Ca(2.87).

3.10.3) What is the Cell Potential of the electrochemical cell in Which
the cell reaction is: Pb2+ + Cd → Pb + Cd2+ ; Given that Eocell = 0.277
volts, temperature = 25oC, [Cd2+] = 0.02M, and [Pb2+] = 0.2M.

Solution
Since the temperature is equal to 25oC, the Nernst equation can be written as follows; Ecell =
E0cell – (0.0592/n) log10Q
Here, two moles of electrons are transferred in the reaction. Therefore, n = 2. The reaction
quotient (Q) is given by [Cd2+]/[Pb2+] = (0.02M)/(0.2M) = 0.1.
The equation can now be rewritten as:
Ecell = 0.277 – (0.0592/2) × log10(0.1) = 0.277 – (0.0296)(-1) = 0.3066 Volts
Thus, the cell potential of this electrochemial cell at a temperature of 25oC is 0.3066 volts.

3.10.4) The Cu2+ ion concentration in a copper-silver electrochemical
cell is 0.1M. If Eo(Ag+/Ag) = 0.8V, Eo(Cu2+/Cu) = 0.34V, and Cell
potential (at 25oC) = 0.422V, find the silver ion concentration.

Solution
Here, the silver electrode acts as a cathode whereas the copper electrode serves as the anode.
This is because the standard electrode potential of the silver electrode is greater than that of
the copper electrode. The standard electrode potential of the cell can now be calculated, as
shown below.
Eocell = Eocathode – Eoanode = 0.8V – 0.34V = 0.46V
Since the charge on the copper ion is +2 and the charge on the silver ion is +1, the balanced cell
reaction is:
2Ag+ + Cu → 2Ag + Cu2+
Since two electrons are transferred in the cell reaction, n = 2. Now, the Nernst equation for this
electrochemical cell can be written as follows.
Ecell = E0cell – (0.0592/2) × log(0.1/[Ag+]2)
0.422V = 0.46 – 0.0296 × (-1 – 2log[Ag+])
Therefore, -2log[Ag+] = 1.283 + 1 = 2.283
Or, log[Ag+] = -1.141
[Ag+] = antilog(-1.141) = 0.0722 M

3.11) Interesting Facts:

• A Voltmeter can’t be used for the precise measurements of galvanic cell….because a
part of the current is drawn by the voltmeter itself, thereby giving lower value than
actual emf.

• A Salt bridge made of KCl solution cannot be used for a cell made of Ag and Pb half
cells because Cl- ions from KCl bridge react with Ag+ ions and Pb+ ions to form AgCl
and PbCl2 respectively. Thus the concentrations of both the half cells will change.

• Only AC currents can be used and not DC in conductance estimation…because when
we use DC current, the products of the electrolysis gets deposited on the electrodes
and set up a back emf thereby increasing the resistance of the electrolyte.

• It is not possible to measure the voltage of an isolated reduction half reaction
because a half cell is incapable for working independently.We can only determine
the relative value of the electrode potential w.r.t any other reference electrodes.

• A dry cell becomes dead after a long time, even if it has not been used because
when the cell is not been used …the acidic NH4Cl slowly corrodes the zinc container
of the dry cell. Thus, dry cell becomes dead after a long time.

3.12) Summary

Electrochemistry is a branch of physical chemistry concerned with the interconversion of chemical and
electrical energy. The electrode potential is the potential difference developed between metal and the
solution. Based on the nature of the metal electrode to gain or lose electrons there exits two types of
electrode potential: oxidation electrode potential and reduction electrode potential of Reference
electrode is an electrode whose potential is arbitrarily taken as zero or is exactly known. We have learnt
various kinds of reference electrodes. These electrodes show slight or no variation of temperature or
any physical factors .Electrochemical cell is a cell that converts electrical energy to chemical energy of
vice-versa.
There 4 kinds: Voltaic(galvanic cell),electrolytic cell, concentration cell and fuel cell. Voltaic cell converts
chemical energy to electrical energy. Electrolytic cell is a non-spontaneous cell since it converts
electrical energy into chemical energy ,the surrounding does work on such kind of cell. Concentration
cell is cell that works when two same metals are dipped in it's own solution differing in concentration. It
is a spontaneous reaction since it converts electrical energy into chemical energy. Fuel cell is based on
the fact that the combustion reactions are redox reactions and hence they can be used to produce
electricity. It is a galvanic cell in which the energy of combustion of fuel is directly converted into
electrical energy called fuel cells. The series is defined as the arrangement of electrodes (metal or non-
metal in contact with their ions) with the electrode half reactions in order of decreasing standard
potential. Fluorine is the strongest oxidizing agent whereas lithium is the strongest reducing agent . This
series is really helpful in finding the relative strenghts of oxidizing and reducing agents. The cell
potential and electrode potential depends on temperature, concentration of solutes and partial
pressures of gases.

This dependence of potential is given by Nernst equation.

3.13) Questionnaire:

3.13.Q1] Very Short Answer:

1. What is galvanic or voltaic cell?

2. What is a half cell (electrode)?

3. What is standard electrode potential?

4. What is the potential of S.H.E?

5. Write name and polarity of an electrode where oxidation takes place in galvanic cell.

6. Name any two metals that can be used in cathodic protection of iron.
7. Articles of iron are generally coated with zinc. Explain.

8. An electrochemical cell stops working after sometime. Explain.

3.13.Q2]Short Answer:

1. What is electrochemical series?
2. What is conductivity cell?
3. Which type of reaction occurs at anode and cathode during electrolysis?
4. Write the electrode reactions taking place in a fuel cell.
5. Give name and formula of compound formed during rusting of iron.
6. Explain why fluorine is the strongest oxidising agent?
7. Lithium metal is the strongest reducing agent. Why?
8. What are the advantages of fuel cell?
9. Why cannot oxidation occur without reduction?
10. What is the difference between cell potential and standard cell potential?
11. Write the cell reaction if the Nernst equation is given by the relation

Ecell=Eocell−RT2Fln[Pb2+][H+]2
12. What do you mean by standard hydrogen electrode (S.H.E)?

3.13Q3]Long Answer:

1. What are fuel cells? Give example.
2. What is corrosion?
3. Give example of a corrosion which is useful?
4. How is potential of an electrode determined?
5. Zinc can reduce hydrogen ion but copper cannot. Why?
6. What is the role of platinum or gold as electrode?

7. Mention the application of electrochemical cells. following galvanic
8. Write Nernst equation and explain the terms involved.
9. Mention some of the applications of electrochemical series.
10. What is electrolytic cell?

11. Corrosion is an electrochemical phenomenon. Explain.
12. Write half cell reactions and balanced chemical equations for the
cells?

(a) Zn(s)+|Zn2+(aq)||Cr3+(aq)|Cr(s)

(b) Pb(s),PbSO4(s)|HSO−4(aq),
H+(aq)||H+(aq),HSO−4(aq)|PbO2(s), PbSO4(s)

(c) Mg(s)|Mg2+(aq)||Sn2+(aq)|Sn(s)

13. The following chemical reaction is occurring in an electrochemical cell:
Mg(s)+2Ag+(0.0001M)→Mg2+(0.10M)+2Ag(s)
Given
EoMg2+/Mg=−2.36V;
EoAg+/Ag=+0.81V

For the cell, calculate/write: E? value for the electrode
2Ag+/2Ag

Standard cell potential (E?) Cell potential (E) Give the symbolic representation of the above cell
Will the above cell reaction be spontaneous?

3.13.Q4] Unsolved Problems:

1. Calculate the electrode potential of copper, if the concentration of CuSO4 is 0.206 at
23.1˚C. Given that E˚Co2+/Cu=+0.34V.

(0.31984V)

2. Calculate the concentration of NiCl2 in the nickel electrode having a potential of
0.16942 V at 24.9˚C. Given that E˚Ni+2/Ni = -0.14V.

(0.1010)

3. Calculate the standard electrode potential of lead electrode, if the electrode potential is
-0.18025 V at 301K and a concentration of Pb2+ solution is 0.0096 M.

(0.1200V)

4. Write the cell reaction and calculate the standard emf of the cell

Cd|Cd+2(1M) || Ag+(1M) |Ag+

E0Cd = -0.403V and E0Ag = 0.799V

(E0cell=1.202V)

5. Construct a cell containing Zn+2 |Zn half cell and hydrogen gas electrode.
What will be the emf of the cell at 250C if [Zn+2] = 0.24M. [H+]=1.6M and

PH2 =1.8 atm? E0Zn = -0.763V
(Ecell=0.786V)

6 . Calculate E0cell and Ecell for the following reaction at 250C
Mg(s) +Sn+2(0.03M) → Mg+2(0.004M) +Sn(s)
E0Mg =-2.37V and E0Sn=-0.14V

(E0cell=2.23V,Ecell=2.226V)
8. Calculate the emf of the cell
Pt|H2(1g,1atm) |H+(0.5M)||KCl(1M)|Hg2Cl2(s)|Hg(l)|Pt at 250C Ecalomel=0.28V.

(Ecell = 0.2978V)
9. Calculate the emf of the cell
Zn(s)|Zn+2(0.008M)||Cr+3(0.01M)|Cr at 250C, E0Zn=-0.763V, E0Cr=-0.74V

(Ecell=0.0456V)

10. Calculate equilibrium constant for the reaction,
Ni(s) + 2Ag+(aq) → Ni+2(aq) + 2Ag(s) at 250C

E0Ni =-0.25V and E0Ag = 0.799V
(K=2.754 x 1035)

11. Consider the cell, Pt|H2(g)|H+(aq)||I-(aq)|I2(s). If the standard cell potential is 0.54V then
the standard potential for cathode half reaction will be?
(+0.54V)

12. Set up the cell consisting of H+(aq)|H2(g) and Pb+2(aq)|Pb(s) electrodes. Calculate the emf

at 250C of the cell if [Pb+2]=0.1M ,[H+]=0.5M and hydrogen gas is at 2atm pressure. E0Pb=-

0.126V (0.1289V)

13. Calculate the potential of the following cell at 250;

Zn|Zn+2(0.6M)||H+(1.2M)|H2(g,1atm)|Pt

E0Zn=-0.763V (0.774V)

References:

https://d1uvxqwmcz8fl1.cloudfront.net/tes/resources/7168639/f443f619-7577-482e-92f5-
15af04020437/image?width=1000&height=190&version=1439946087934
https://i.ytimg.com/vi/FnJ0V7B7nKo/maxresdefault.jpg
https://www.corrosionpedia.com/definiti/970/reference-electrodeon
https://byjus.com/chemistry/standard-hydrogen-electrode/
https://eduladder.com/viewquestions/3618/Discuss-the-construction-and-working-
ofcalomelelectrode
chromeextension://oemmndcbldboiebfnladdacbdfmadadm/http://www.griet.ac.in/nodes/EC_
UNIT_2.p df https://pubs.rsc.org/en/content/articlelanding/2015/AY/C5AY01724F#!divAbstract
https://www.ncbi.nlm.nih.gov/pubmed/15352497 https://www.askiitians.com/iit-

jeechemistry/physical-chemistry/electrode-potential.aspx

France, Colin (2008), The Reactivity Series of Metals.

NCERT, A textbook for Chemistry http://wiki.engageeducation.org.au/chemistry/unit-4/area-of-
study-2-supplying-and-
usingenergy/electrochemical- series/
(From a reference book of Abhijit Mallick)
http://www.indicareer.com/entranceexams/mhtcet/chemistry/Electrochemistry-2.html

4.CORROSION

4.1 INTRODUCTION

Corrosion is one of the most significant problems faced by advanced industrial societies. It is
estimated that about 30 to 40% of iron and steel produced annually is used just to replace
the rusted iron materials. Corrosion costs of automobiles- fuel systems, radiators, exhaust
systems and bodies- are in the billions. Corrosion touches all appliances used inside and
outside home, on the road, on the sea, in plants and in aerospace vehicles. Total annual
cause of floods, hurricane, fires, lightning and earthquakes are less than the costs incurred
due to corrosion.

Metals are thermodynamically unstable and found abundantly in nature coma so that they
require large amounts of energy to get converted into useful Engineering Materials. Thus,
the metals acquiring more energy have a greater tendency to corrode, whereas those
metals requiring less energy have a lower tendency to corrode. Examples of the former
include magnesium zinc aluminium and steel; whereas examples of the latter include gold,
silver and Platinum. Corrosion is spontaneous slow chemical interaction of metal or alloy
with its environment coma resulting in the formation of one of its compounds such as oxide,
hydrated oxide, carbonate, sulphide sulphate, etc. Extraction of metals from their ores is an
endothermic process. Pure metals, being highly energetic, have natural tendency to revert
back to their combined States.

The term corrosion is sometimes also applied to the degradation of plastics concrete and
wood, but generally refers to metals. The most familiar example of corrosion is rusting of
iron exposed to the atmosphere conditions.

4.1.2 HISTORICAL OVERVIEW OF CORROSION

(Pliny) (Robert Bolye)

The word corrosion is as old as the earth, but it has been known by different names. A
Roman philosopher, Pliny (AD 23–79) wrote about the destruction of iron in his essay
‘Ferrum Corrumpitar’. There is a historical record of observation of corrosion by several
writers, philosophers and scientists, but there was little curiosity regarding the causes and
mechanism of corrosion until Robert Boyle wrote his ‘Mechanical Origin of Corrosiveness.’
Philosophers, writers and scientists observed corrosion and mentioned it in their writings:

 Pliny the elder (AD 23–79) wrote about spoiled iron.
 Herodotus (fifth century BC) suggested the use of tin for protection of iron.
 Lomonosov (1743–1756).
 Austin (1788) noticed that neutral water becomes alkaline when it acts on iron.
 Thenard (1819) suggested that corrosion is an electrochemical phenomenon.
 Hall (1829) established that iron does not rust in the absence of oxygen.
 Davy (1824) proposed a method for sacrificial protection of iron by zinc.
 De la Rive (1830) suggested the existence of microcells on the surface of zinc.

4.2 Definition of corrosion

Corrosion is the deterioration or destruction of metals and alloys in the presence of an
environment by chemical or electrochemical means. In simple terminology, corrosion
processes involve reaction of metals with environmental species.

As per IUPAC, “Corrosion is an irreversible interfacial reaction of a material (metal, ceramic,
polymer) with its environment which results in its consumption or dissolution into the
material of a component of the environment. Often, but not necessarily, corrosion results in
effects detrimental to the usage of the material considered. Exclusively physical or
mechanical processes such as melting and evaporation, abrasion or mechanical fracture are
not included in the term corrosion.”

With the knowledge of the role of various microorganisms present in soil and water bodies,
the definition for corrosion need be further widened to include microbiallyinfluenced
factors.

4.2.1 Classification of Corrosion

Corrosion can be classified in different ways, such as:
 Dry or Chemical corrosion
 Wet or Electrochemical corrosion

4.3 Dry Corrosion (Chemical corrosion)

Dry corrosion is a form of corrosion that does not require the presence of a liquid
electrolyte. It is also called as atmospheric corrosion. Sometimes, this type of damage is
called dry corrosion or scaling. The term oxidation is ambivalent because it can either refer
to the formation of oxides or to the mechanism of oxidation of a metal (i.e., its change to a
higher valence than the metallic state).
Dry corrosion occurs when there is no moisture or water to aid corrosion. Dry corrosion is
taken place in the presence of gases in the air.
This type of corrosion is due to the direct chemical attack of metal surfaces by the
atmospheric gases such as oxygen, halogen, hydrogen sulphide, sulphur dioxide, nitrogen or
anhydrous inorganic liquid, etc.
The chemical corrosion is defined as the direct chemical attack of metals by the atmospheric
gases present in the environment.
Example: (i) Silver materials undergo chemical corrosion by Atmospheric H2S gas.

(ii) Iron metal undergo chemical corrosion by HCl gas.

4.3.2 Classifications of Dry Corrosion:

Dry Corrosion

Oxidation Hydrogen Liquid metal
Corrosion Embrittlement corrosion

Corrosion by Corrosion by
oxygen other gasses

4.3.3 Oxidation Corrosion:

1) Corrosion by Oxidation: -

It takes place by direct reaction of oxygen on metal. It occurs in absence of moisture . It
mostly occurs at ordinary temperature.

Mechanism: When a metal is exposed to air it gets oxidized by loosing its valence electrons
& reduction of oxygen take place

M + n/4 (O2)  Mn+ + O2-

At point of contact of Mn+ & O2- metallic oxide will form & that metallic oxide scale forms a
barrier to restrict further oxidation of inside the metal. Since size of cation (Mn+) is smaller
than the anion, hence the cation will diffuse much faster than the anion through the scale
for continuation of oxidation, it can be possible if the metallic oxide barrier is sufficiently
porous. The nature of oxide film plays very important role in oxidation corrosion.

i.) When oxide film is stable and tightly adhering, it will act as protective coating
and corrosion is prevented.

ii.) When oxide film is unstable and has tendency to decompose back to metal and
oxygen, it does not go into oxidation corrosion.

iii.) When film is volatile then metal surface again come into contact with air and
oxidation take place.

iv.) If film is sufficiently porous then continuous oxidation takes place.

These cases are listed below into the following types:

(i) Stable layer: A Stable layer is fine grained in structure and can get adhered tightly

to the parent metal surface. Hence, such layer can be of impervious nature (i.e. which cuts-

off penetration of attaching oxygen to the underlying metal). Such a film behaves as

protective coating in nature, thereby shielding the metal surface. The oxide films on Al, Sn,

Pb, Cu, Pt, etc., are stable, tightly adhering and impervious in nature.

(ii) Unstable oxide layer: This is formed on the surface of noble metals such as Ag, Au,
Pt. As the metallic state is more stable than oxide, it decomposes back into the metal and
oxygen. Hence, oxidation corrosion is not possible with noble metals.

(iii) Volatile oxide layer: The oxide layer film volatilizes as soon as it is formed. Hence,
always a fresh metal surface is available for further attack. This causes continuous corrosion.
MoO3 is volatile in nature.

(iv) Porous layer: The layer having pores or cracks. In such a case, the atmospheric
oxygen have access to the underlying surface of metal, through the pores or cracks of the
layer, thereby the corrosion continues unobstructed, till the entire metal is completely
converted into its oxide.

2) Corrosion by other gases: -

Some gases react with certain metal and forms a protective or non-protective layer on the
metallic surface . Due to chemical combination of metals with gases, metals undergo
corrosion. The extend of corrosion depends upon:
i.)The Nature of the environment.
ii.) Nature of oxide film formed.

4.3.4 Liquid metal corrosion: -

When a liquid metal is allowed to flow over solid metal at high temperature is called liquid
metal corrosion. Due to this solid metal gets weak. Example- In nuclear reactor sodium
metal is used as coolant & it leads to corrosion of cadmium.

4.3.5 Hydrogen Embrittlement:

Loss in ductility of a material in the presence of hydrogen is known as hydrogen
embrittlement.
Mechanism: This type of corrosion occurs when a metal is exposed to hydrogen
environment. Iron liberates atomic hydrogen with hydrogen sulphide in the following way.

Fe + H2S → FeS + 2H
Hydrogen diffuses into the metal matrix in this atomic form and gets collected in the voids
present inside the metal. Further, diffusion of atomic hydrogen makes them combine with
each other and forms hydrogen gas.

H + H → H2↑

Collection of these gases in the voids develops very high pressure, causing cracking or
blistering of metal.

Decarburisation:
The presence of carbon in steel gives sufficient strength to it. But when steel is exposed to
hydrogen environment at high temperature, atomic hydrogen is formed.

H2  2H

Atomic hydrogen reacts with the carbon of the steel and produces methane gas.

C + 4H → CH4

Hence, the carbon content in steel is decreases. The process of decrease in carbon content
in steel is known as decarburization. Collection of methane gas in the voids of steel develops
high pressure, which causes cracking. Thus, steel loses its strength.

Lets test your Knowledge

1. Dry corrosion is also called as_________

A] Chemical corrosion B] Electrochemical corrosion

C] Wet corrosion C] Oxidation corrosion

2. Anhydrous inorganic liquid metal surface in absence of moisture undergoes ___________

A] Wet corrosion B] Dry corrosion
C] Galvanic corrosion D] Pitting corrosion

3. Corrosion is uniform in__________
A]Dry corrosion B]Wet corrosion
C] Pitting corrosion D] Water line corrosion

4. Dry corrosion takes place in__________ B] Heterogeneous solutions
A] Homogeneous solutions D] Both homogeneous and heterogeneous
C]Neither homogeneous nor heterogeneous
solutions

5. Volatile oxidation corrosion product of a metal is
A] Fe2O3 B] MoO3
C] Fe3O4 D] FeO

6. Hydrogen embrittlement happens in which environment:

A] Humid B] High pressure

C] High Temperature D] Cold environment

Anwers

1 2 3456

A BADBC

4.4 Wet Corrosion (Electrochemical corrosion)

This type of corrosion occurs: (i) where a conducting liquid is in contact with metal (ii)
when two dissimilar metals or alloys are either immersed or dipped partially in a solution.
This corrosion occurs due to the existence of separate ‘anodic’ and ‘cathodic’ areas/parts
between which current flows through the conducting solution. At anodic area, oxidation
reaction (i.e. liberation of free electrons) takes place, so anodic metal is destroyed by either
dissolving or assuming combined state (such as oxide, etc). Hence corrosion always occurs
at anodic areas.