1587884932985_final book 26-04-20

ALCHEMY
AN ENGINEERING
CHEMISTRY E-BOOK

PREFACE

Alchemy, an Engineering Chemistry e-book has been developed by First Year
Engineering students of Vidyalankar Institute of Technology under the guidance of
Prof. Sonaali Borkar, Assistant Professor, Engineering Chemistry.

The objective of writing this book is to provide a smooth, interactive and easy
approach to learning Engineering Chemistry. The book is full of visual
presentations taking into consideration the needs of the millennial learners.
Information, pictures, diagrams, videos and animations make this book engaging,
simple to understand and motivate self-learning.

The contents of the book are taken from various online resources, reference books
on Engineering Chemistry and teachers notes.

This book is not developed for any commercial purpose but aims to serve as a sort
of educational guide for all the First Year Engineering students of VIT. Any other
student studying Engineering Chemistry is also welcome to refer to this ebook.

This is our first attempt so kindly give your feedback and valuable inputs to
improve the content via email to [email protected]

Pictures of all members and basic info

Prof. Sonaali Borkar,
Assistant Professor,
Engineering Chemistry, VIT

Parth Yadav Dishant
Computer Engineering, VIT Computer Engineering, VIT

Kaveya Sivaprakasam Nimish Samant
Computer Engineering, VIT Computer Engineering, VIT

Karthik Subhash Ritika Lath
Computer Engineering, VIT Computer Engineering, VIT

Shreyas Kunturkar Govind Mundhe
Computer Engineering, VIT Electronics &Telecommunication
Engineering, VIT

Tuba Momin Ayushi Bhagat
Computer Engineering, VIT Computer Engineering, VIT

Course Course Name Teaching Scheme Credits Assigned
Code Engineering (Contact Hours) Theory Tut. Pract. Total
FEC203 Chemistry-II
Theory Pract. Tut.
Course Course Name
Code 2 - - 2 -- 2
Engineering
FEC203 Chemistry-II Examination Scheme

Theory

Internal Assessment End Exam. Term Pract. Total
Test1 Test 2 Avg. Sem. Duration Work /oral
Exam. (in Hrs)

15 15 15 60 2 -- -- 75

Objectives
The concepts developed in this course will aid in quantification as well as understand the
applications of several concepts in Chemistry that have been introduced at the 10 + 2
levels in schools.

Outcomes: Learners will be able to…
1. Distinguish the ranges of the electromagnetic spectrum used for exciting different
molecular energy levels in various spectroscopic techniques.
2. Illustrate the concept of emission spectroscopy and describe the phenomena of
fluorescence and phosphorescence in relation to it.
3. Explain the concept of electrode potential and nernst theory and relate it to electrochemical
cells.
4. Identify different types of corrosion and suggest control measures in industries.
5. Illustrate the principles of green chemistry and study environmental impact.
6. Explain the knowledge of determining the quality of fuel and quantify the oxygen required
for combustion of fuel.

Module Detailed Contents Hrs.

Principles of Spectroscopy:

Introduction: Principle of spectroscopy, Definition,Origin of spectrum,

01 Classification of spectroscopy – atomic and molecular, selection rules. 02

Table of relation between electromagnetic spectrum, types of spectroscopy and

energy changes.

Applications of Spectroscopy

Emission spectroscopy- Principle, Instrumentation and applications ( Flame

02 Photometry) 04

Introduction to florescence and phosphorescence, Jablonski diagram, application

of fluorescence in medicine only.

Concept of Electrochemistry
Introduction, concept of electrode potential, Nernst equation, types of
03 electrochemical cells, concept of standard electrode with examples, 02
electrochemical series, simplenumericals.

University of Mumbai, First Year Engineering, (Common for all Branches of Engineering) REV2019 ‘C’ Scheme 40/61

Corrosion:

Definition, Mechanism of Corrosion- (I) Dry or Chemical Corrosion-i) Due to

oxygen ii)Due to other gases.

(II)Wet or Electrochemical corrosion- Mechanism

i) Evolution of hydrogen type ii) Absorption of oxygen.

Types of Corrosion- Galvanic cell corrosion, Concentration cell corrosion

04 (differential aeration principle), Pitting corrosion, Intergranular corrosion, Stress 06
corrosion.

Factors affecting the rate of corrosion- (i)Nature of metal, (ii)Nature of corroding

environment.

Methods of corrosion control- (I)Material selection and proper designing,(II)

Cathodic protection- i) Sacrificial anodic protection ii) Impressed current

method,(III) Metallic coatings- only Cathodic coating (tinning) and anodic

coatings (Galvanising)

Green Chemistry and Synthesis of drugs
Introduction – Definition, significance

05 Twelve Principles of Green chemistry, numerical on atom economy, 04
Conventional and green synthesis of Adipic acid, Indigo, Carbaryl, Ibuprofen,

Benzimidazole, Benzyl alcohol, % atom economy and their numericals.

Green fuel- Biodiesel.

Fuels and Combustion

Definition, classification, characteristics of a good fuel, units of heat (no

conversions).

Calorific value- Definition, Gross or Higher calorific value & Net or lower
calorific value, Dulong’s formula & numerical for calculations of Gross and Net

06 calorific values. 06
Solid fuels- Analysis of coal- Proximate and Ultimate Analysis- numerical

problems and significance.

Liquid fuels- Petrol- Knocking, Octane number, Cetane number, Antiknocking

agents, unleaded petrol, oxygenates (MTBE), catalytic converter.

Combustion- Calculations for requirement of only oxygen and air (by weight and

by volume only) for given solid & gaseous fuels.

Assessment

Internal Assessment Test

Assessment consists of two class tests of 15 marks each. The first class test is to be conducted
when approx. 40% syllabus is completed and second class test when additional 35% syllabus is
completed. Duration of each test shall be one hour.

End Semester Examination

In question paper weightage of each module will be proportional to number of respective
lecture hours as mention in the syllabus.

1. Question paper will comprise of 6 questions, each carrying 15 marks.
2. Question number 1 will be compulsory and based on maximum contents of the syllabus
3. Remaining questions will be mixed in nature (for example, if Q.2 has part (a) from module

3 then part (b) will be from other than module 3)
4. Total four questions need to be solved.

University of Mumbai, First Year Engineering, (Common for all Branches of Engineering) REV2019 ‘C’ Scheme 41/61

Recommended Books :

1. Engineering Chemistry - Jain & Jain, DhanpatRai
2. Engineering Chemistry – Dara & Dara, S Chand
3. Green Chemistry: A textbook – V.K.Ahluwalia, Alpha Science International
4. Fundamentals of Molecular Spectroscopy ( 4th Edition) - C.N.Banwell, Elaine M.

McCash,
Tata McGraw Hill.
5. Elementary Organic Spectroscopy- Y.R.Sharma, S.Chand and Co.
6. A Text Book of Engineering Chemistry - ShashiChawla, DhanpatRai
7. Engineering Chemistry – Payal Joshi &Shashank Deep (Oxford University Press)

University of Mumbai, First Year Engineering, (Common for all Branches of Engineering) REV2019 ‘C’ Scheme 42/61

TABLE OF CONTENTS

1. Principles of Spectroscopy

1.1 Introduction

1.1.1 Origin of Spectrum

1.2 Electromagnetic Radiations
1.3 Electromagnetic Spectrum
1.4 Classification of Spectroscopy

1.4.1 Atomic Spectroscopy
1.4.2 Molecular Spectroscopy

1.4.2.1 Types of Spectroscopy

1.4.3 Selection Rules for Electronic Transitions
1.4.4 Types of Molecular Energies

Let us Revise

1.5 Summary

1.6 Short Answer Questions
1.7 Long Answer Questions

2. Applications of Spectroscopy

2.1 Introduction
2.2 Emission Spectroscopy

2.2.1 Flame Photometry
2.2.1.1 Flame Photometer

2.3 Fluorescence Spectroscopy
2.4 Phosphorescence Spectroscopy

2.4.1 Fluorescence and Phosphorescence
2.4.2 Fluorescence
2.4.3 Phosphorescence
2.4.4 Applications of Fluorescence Spectroscopy

2.5 Jablonski Diagram

2.5.1 Radiative and Non-Radiative Transitions

2.5.1.1 How to Interpret a Jablonski Diagram

Let us Revise

2.6 Summary
2.7 Short Answer Questions
2.8 Long Answer Questions

3. ELECTROCHEMISTRY

3.1 Introduction

3.1.1 Definition
3.1.2 History (From Volta to Faraday)

3.2 Electrode Potential

3.2.1 Types of Electrode Potential
3.2.2 Classification of electrodes

3.3 Concept of Standard Electrode Potential

3.3.1 Reference Electrode
3.3.2 Types of Reference Electrode
3.3.3 Standard Hydrogen Electrode
3.3.4 Calomel Electrode (SHE)

3.4 Electrochemical Cell
3.5 Types of Electrochemical Cell

3.5.1 Voltaic Cell (Galvanic Cell)
3.5.2 Electrolytic Cell
3.5.3 Concentration Cell
3.5.4 Fuel Cell

3.6 Electrochemical Series

3.6.1 Characteristics Of Electrochemical Series
3.6.2 Applications Of Electrochemical Series
3.6.3 Mnemonics For Electrochemical Series

3.7 Nernst Equation

3.7.1 Derivation Of Nernst Equation
3.7.2 Applications Of Nernst Equation
3.7.3 Limitations Of Nernst Equation

3.8 Relation Between Corrosion And Electrochemistry

3.9 Recent Development In Electrochemistry

3.9.1 Paper Based Analytical Devices (PAD’s)
3.9.2 Biosensors

3.10 Solved Examples On Nernst Equation
3.11 Interesting Facts
3.12 Summary
3.13 Questionnaire

4.CORROSION

4.1 Introduction

4.1.1 Historical overview of corrosion

4.2 Definition of corrosion

4.2.1 Classification of corrosion

4.3 Dry corrosion (Chemical corrosion)

4.3.1 Classification of dry corrosion

4.3.2 Oxidation of corrosion

4.3.3 Liquid metal corrosion

4.3.4 Hydrogen Embrittlement

4.4 Wet corrosion (Electrochemical Corrosion)

4.4.1 Hydrogen Evolution Type

4.4.2 Oxygen Absorption Type

4.4.3 Electrochemical Theory of Wet corrosion

4.5 Corrosion Cell

4.6 Corrosion Mechanism of Iron oxide (Rust) Formation

4.7 Difference Between Dry and Wet Corrosion
4.8 Classification of Electrochemical corrosion

4.8.1 Pitting corrosion
4.8.2 Intergranular Corrosion
4.8.3 Stress Corrosion Cracking
4.8.4 Galvanic Corrosion
4.8.5 Concentration Cell Corrosion
4.8.6 Crevice Corrosion
4.8.7 Microbiological Corrosion
4.8.8 Uniform Corrosion
4.8.9 Dealloying
4.8.10 Fretting Corrosion

4.9 Factors that affect the rate of corrosion

4.9.1 Nature of the metal
4.9.2 Nature of the corroding environment

4.10 Methods of corrosion control

4.10.1 Improvement of design
4.10.2 Proper selection Of materials
4.10.3 Cathodic protection
4.10.4 Anodic protection
4.10.5 Modifying the environment

4.11 Use of protective coatings

4.11.1 Inorganic protective coating
4.11.2 Methods of application of metallic coatings

4.12 Characteristics of water provide to the corrosion process
4.13 Worldwide impact of corrosion (Statistics)

4.14 Summary

4.15 Model questions

4.15.1 MCQ type

4.15.2 Short answer type

4.15.3 Long answer type

5. GREEN CHEMISTRY

5.1. Introduction

5.1.1. History
5.1.2. Father of Green Chemistry
5.1.3. Definition
5.1.4. What is Green Chemistry
5.1.5. Goals of Green Chemistry

5.2. 12 Principles of Green Chemistry

5.2.1. Waste Prevention
5.2.2. Atom Economy
5.2.3. Less Hazardous Chemical Synthesis
5.2.4. Designing Safer Chemicals
5.2.5. Safer Solvents and Auxiliaries
5.2.6. Design for Energy Efficiency
5.2.7. Use of Renewable Feedstocks
5.2.8. Reduce Derivatives
5.2.9. Catalyst
5.2.10. Design for Degradation
5.2.11. Real - Time Pollution Prevention
5.2.12. Safer Chemistry for Accident Prevention

5.3. Atom Economy
5.4. Benefits of green chemistry

5.5. Green synthesis of

5.5.1. Adipic Acid
5.5.2. Indigo
5.5.3. Carbaryl
5.5.4. Ibuprofen
5.5.5. Benzimidazole

5.6. Biodiesel

5.6.1. Introduction
5.6.2. Advantages
5.6.3. Disadvantages

5.7. Green computing

5.7.1. Introduction
5.7.2. Origin
5.7.3. Pathways to green chemistry
5.7.4. Importance
5.7.5. Recommendation
5.7.6. Conclusion

5.8. Summary

5.9. Solved Numerical

5.10. Unsolved Numericals
5.11. Short questions
5.12. Long questions

6. FUELS and COMBUSTION

6.1 Introduction

6.1.1 Definition

6.2 Classification of fuels

6.3 Characteristics of good fuel
6.4 Calorific Value

6.4.1 Gross Calorific Value
6.4.2 Net Calorific Value

6.5 Dulong’s Formulae
6.6 Solid Fuel
6.7 Types of solid fuel
6.8 Coal

6.8.1 Properties of Coal
6.8.2 Types of Coal
6.8.3 Analysis of Coal

6.9 Liquid Fuel

6.9.1 Types of liquid fuel
6.9.2 Refining of crude oil

6.10 Knocking

6.10.1 Petrol engine
6.10.2 Diesel engine
6.10.3 Octane number
6.10.4 Cetane number
6.10.5 Anti-knocking agents
6.10.6 Unleaded petrol
6.10.7 Catalytic converter

6.11 Theory for solving problems on combustion
6.12 Summary
6.13 Questions

6.13.1 Short answer the following
6.13.2 Long answer the following
6.13.3 Numericals

1
Principles of Spectroscopy

1.1 Introduction

The term "spectroscopy" defines a large number of techniques that use radiation to obtain
information on the structure and properties of matter. The basic principle shared by all
spectroscopic techniques is to shine a beam of electromagnetic radiation onto a sample,
and observe how it responds to such a stimulus. The response is usually recorded as a
function of radiation wavelength. A plot of the response as a function of wavelength is
referred to as a spectrum.
Molecular Spectroscopy is an experimental process to measure as to which frequencies of
radiation are absorbed or emitted by a particular substance and then attempting to
correlate the patterns of the energy absorption or emission with details of molecular
structure. In the quantum mechanical terms, electromagnetic radiations are considered to
be propagated in discrete packets of energy called photons. These photons have very
specific energies and are quantised. The energy of a photon is derived by the relationship

E=hv
All the energy states of a molecule are quantised. It is this fact that makes spectroscopy
possible.

An atom or a molecule can be made to undergo a transition from energy state E1 to a higher
energy state E2 by irradiating the atom or molecule with electromagnetic radiation
corresponding to the energy difference between states E1 and E2. When the atom or
molecule returns from state E2 to E1, an equivalent amount of energy is emitted. Thus,
excitation of a molecule involves absorption of specific quanta of electromagnetic radiation.
The photon of electromagnetic radiation will be only absorbed by a molecule if its energy
corresponds exactly to an energy difference E between two states of the molecule.

The energy difference between e.g., the excited state (E2) and the ground state (E1) will
correspond to a certain frequency (v) or wavelength (λ) of electromagnetic radiation, which
will depend on the type of transition.

The relation between the energy of a transition and frequency is given by

E= hv

E=hc/λ

In conclusion, therefore spectroscopy is the study of interaction of molecules and
electromagnetic radiations.

1.1.1 Origin of Spectrum

A spectrum is simply a plot of the amount of light which is
absorbed versus the frequency (or wavelength) of the light. The
spectrum yields information about the spacing of the energy
levels of the molecule. Since these energy levels depend on the
structure of the molecules, this information is therefore used to
determine the structure of the compound. Next figure illustrates
in a pictorial fashion the various, rather arbitrary, regions into
which electromagnetic radiation has been divided.

The boundaries between the regions are by no means precise,
although the molecular processes associated with each region are
quite different.

1.2 Electromagnetic Radiations

Spectroscopy may be defined as
the study of the interaction of
electromagnetic waves and
matter. Electromagnetic Radiation,
of which visible light forms an
obvious but very small part, may
be considered as a simple
harmonic wave propagated from a
source and travelling in straight
lines except when reflected or
refracted. EM radiation is created
when an atomic particle, such as
an electron, is accelerated by an electric field, causing it to move. The movement produces
oscillating electric and magnetic fields, which travel at right angles to each other in a bundle
of light energy called a photon. Photons travel in harmonic waves at the fastest speed
possible in the universe: 186,282 miles per second (299,792,458 meters per second) in a
vacuum, also known as the speed of light. The waves have certain characteristics, given as
frequency, wavelength or energy.

A wavelength is the distance between two consecutive peaks of a wave. This distance is
given in meters (m) or fractions thereof. Frequency is the number of waves that form in a
given length of time. It is usually measured as the number of wave cycles per second, or
hertz (Hz). A short wavelength means that the frequency will be higher because one cycle
can pass in a shorter amount of time, according to the University of Wisconsin. Similarly, a
longer wavelength has a lower frequency because each cycle takes longer to complete.

1.3 Electromagnetic Spectrum

EM radiation spans an enormous range of wavelengths and frequencies. This range is known
as the electromagnetic spectrum. The EM spectrum is generally divided into seven regions,
in order of decreasing wavelength and increasing energy and frequency.

Radio waves are at the lowest range of the EM spectrum, with frequencies of up to about
30 billion hertz, or 30 gigahertz (GHz), and wavelengths greater than about 10 millimetres
(0.4 inches). Radio is used primarily for communications including voice, data and
entertainment media.

Microwaves fall in the range of the EM spectrum between radio and IR. They have
frequencies from about 3 GHz up to about 30 trillion hertz, or 30 terahertz (THz), and
wavelengths of about 10 mm (0.4 inches) to 100 micrometres (μm), or 0.004 inches.
Microwaves are used for high-bandwidth communications, radar and as a heat source for
microwave ovens and industrial applications.

Infrared is in the range of the EM spectrum between microwaves and visible light. IR has
frequencies from about 30 THz up to about 400 THz and wavelengths of about 100 μm
(0.004 inches) to 740 nanometres (nm), or 0.00003 inches. IR light is invisible to human
eyes, but we can feel it as heat if the intensity is sufficient.

Visible light is found in the middle of the EM spectrum, between IR and UV. It has
frequencies of about 400 THz to 800 THz and wavelengths of about 740 nm (0.00003 inches)
to 380 nm (.000015 inches). More generally, visible light is defined as the wavelengths that
are visible to most human eyes.

Ultraviolet light is in the range of the EM spectrum between visible light and X-rays. It has
frequencies of about 8 × 1014 to 3 × 1016 Hz and wavelengths of about 380 nm (.000015
inches) to about 10 nm (0.0000004 inches). UV light is a component of sunlight; however, it
is invisible to the human eye. It has numerous medical and industrial applications, but it can
damage living tissue.

X-Rays are roughly classified into two types: soft X-rays and hard X-rays. Soft X-rays
comprise the range of the EM spectrum between UV and gamma rays. Soft X-rays have
frequencies of about 3 × 1016 to about 1018 Hz and wavelengths of about 10 nm (4 ×
10−7 inches) to about 100 picometers (pm), or 4 × 10−8 inches. Hard X-rays occupy the same
region of the EM spectrum as gamma rays. The only difference between them is their
source: X-rays are produced by accelerating electrons, while gamma rays are produced by
atomic nuclei.

Gamma-rays are in the range of the spectrum above soft X-rays. Gamma-rays have
frequencies greater than about 1018 Hz and wavelengths of less than 100 pm (4 ×
10−9 inches). Gamma radiation causes damage to living tissue, which makes it useful for
killing cancer cells when applied in carefully measured doses to small regions. Uncontrolled
exposure, though, is extremely dangerous to humans.

Let us Revise:
1. Which electromagnetic radiation has the lowest energy?
2. Which electromagnetic Radiation is used by mobile phones?
3. How are EM Waves generated?
4. What is the use of Radio waves?
5. Which type of UV has the highest hazard level?

1.4 Classification of Spectroscopy

1.4.1 Atomic Spectroscopy

The Spectroscopy of Hydrogen: Bohr’s Model.

Hydrogen atoms can be excited with an electrical discharge in vessel containing the gas at
low pressure. In the visible the line spectrum is the famous converging series. In 1885 the
Swiss school teacher Balmer showed that the spectrum was well described by the
Geometric Series.

where G is an empirical constant. In 1889, Rydberg suggested that the Balmer formula was a
special case of the expression

where n1=2. The redefined formula RH has come to be known as the Rydberg constant. In
1906 and 1908 Lyman and Pashcen confirmed this expression by finding additional series (in
the ultraviolet and infrared) described by the same expression with n1=1 or n1=3.
In 1913 Bohr invented a simple model of atom quantization. If the angular momentum of
circular orbits was fixed in units of h, the energy of the orbits would also be fixed and
indexed by a quantum number n, according to

In a transition between states n1 and n2, the energy difference would be balanced by
emission or absorption of a photon with wavelength given by the Planck formula

so, we have derived from first principles the Rydberg equation and the Rydberg constant.

The modern value of RH is given in order to convey the kind of precision that is typical of the
Atomic Spectroscopy. Note that, following convention since the days of early spectroscopy,
the units of Rydberg are in terms of the wave number or inverse length.
The value of RH given above is actually the value that would be appropriate if the nucleus
were fixed or equivalently, of infinite mass. In practice, the mass of electron should be
replaced by the reduced mass, m/(m +MP), where MP is the mass of the nucleus, in this, of
the proton. Thus, the spectrum has slight dependence on the mass of the nucleus. It is easy
to show that the frequency shift between hydrogen and deuterium is

where we approximate MD by 2MP. Thus, we can measure m/MP by comparing the spectra
of hydrogen and deuterium.
Bohr’s derivation of the Rydberg’s formula using quantized atomic orbits is one of the
landmarks of physical science. Bohr’s is a classic calculation using simple mathematics. You
should follow the treatment in your references and learn to do it yourself. We now know
that the more sophisticated Schrodinger’s treatment, solving for the electron wavefunction
in the spherically symmetric Coulomb potential gets same energy eigen-values. In
retrospect, Bohr’s simple model is not much more than dimensional analysis built around
the quantum hypothesis. That fact that it works contains a deep lesson about the
fundamental simplicity and solvability of nature In the Schrodinger solution, we find that
each electronic state is specified by three quantum numbers, n, l, and m. the principle
quantum number n specifies the radial form of the wavefunction and energy, and is logically
similar to index for the quantized Bohr’s orbits. The quantum number l is restricted to the
range l ≤n-1, and specifies the angular part of the wavefunction and the total angular
momentum for the state: |L|=√l(l+1)h/2. The magnetic quantum number, m, specifies the
projection of the total angular momentum on a given axis:

In the absence of an external magnetic field, we do not expect the energy of the state to
depend on its orientation in space; states with different values of m, therefore have the
same energy and are said to be degenerate. What is perhaps a bit surprising is that the
energy is also independent of the angular momentum, l. This is an accidental consequence
of the spherically symmetric, 1/r form of the Coulomb potential. The energy levels of
hydrogen depend only on n, and all the states of different l and m for a given n are
degenerate.

Note that the energy only depends on the index n. note that the transitions as specified
always invoke a change of 1 unit of angular momentum l. Because the states are
degenerate, this selection rule does not show up in the Radiation Spectrum of Hydrogen,
but it will be of consequence elsewhere.

In the experiment on Atomic Spectroscopy the Balmer series in hydrogen will be measure
with a modern industrial-grade spectrometer, utilizing fibre optics, precision grating, and a
CCD based readout is used to measure the Balmer Spectrum of Hydrogen.

1.4.2 Molecular Spectroscopy

This far we have been mainly confirmed with the emission from single atoms, i.e., spectra
characterized by sharp lines arising from the emission of photons of particular energies
when an electron makes a transition from an excited atomic state to a lower energy state.
We will now look at the spectrum of a simple diatomic molecular system, N2.
A cursory inspection of the system of molecular nitrogen reveal not isolated sharp lines but
broad bands of emission frequencies. This more complicated spectrum is the result of
excitation of the additional degrees of freedom that molecules exhibit compared to single,
namely the vibrational and rotational motion of the excited molecule. The emission due to
electronic transitions is now convoluted with the special features associated with vibrational
and rotational excitations.
The details of this arrangement are explained in the references, particularly Eisberg and
Resnick, Herzberg and Banwell; however, it is easy to extract some of the overall features of
the spectrum for an intuitive understanding. Because the energies associated with
vibrational and rotational degrees of freedom are quite different, we can easily distinguish
these two types of excitation from each other in the spectrum. This separation is embodied
more formally as the Born-Oppenheimer approximation. In this case we can treat the
vibrational and rotational features of the spectrum as separate contributions superimposed
on the overall electronic transition between molecular states:
Let’s start with the vibration of the molecule. A useful model for vibrational excitations is
the quantum harmonic oscillator which has energy levels given by:

Where n is the vibrational quantum number, an integer o corresponds to the fundamental
frequency of the simple harmonic oscillator (SHO). Transitions in which n = -1 lead to an
emission spectrum of equally spaced lines. Coming now to the rotational excitations, if you
zoom in in one of the peaks in the comb of vibrational lines you will see some fine structure.

This fine structure is replicated on the top of each vibrational lines. It arises from optical
emission associated with transitions between the quantized energy levels of a rigid rotator:

Where J=0, 1, 2,………, h is the Planck’s constant, and I is the moment of inertia of the N2
molecule. Note that the energy levels of the rotational states given by this equation are not
equally spaced.

1.4.2.1 Types of Spectroscopy

1. Ultraviolet and Visible Spectroscopy: Ultraviolet (UV) and visible (Vis) spectroscopy
analyses compounds using the electromagnetic radiation spectrum from 10 nm to
700 nm. Many atoms are able to emit or absorb visible light, and it is this absorption
or reflectance that gives the apparent colour of the chemicals being analysed.
Wavelengths of light all have a particular energy associated with them, and it is only
light with the right amount of energy that causes transitions from one level to
another for absorption. For larger gaps between energy levels, more energy is
required for promotion to the higher energy level, so there will be higher frequency
and shorter wavelength absorbed.

2. Infrared Spectroscopy: Infrared (IR) analyses compounds using the infrared
spectrum, which can be split into near IR, mid IR and far IR. Near IR has the greatest
energy and can penetrate a sample much deeper than mid or far IR, but due to this it
is also the least sensitive. Infrared spectroscopy is not as sensitive as UV/Vis
spectroscopy due to the energies involved in the vibration of atoms being smaller
than the energies of the transitions. IR uses the principle that molecules vibrate, with
bonds stretching and bending, when they absorb infrared radiation. IR spectroscopy
works by passing a beam of IR light through a sample, and for an IR detectable
transition, the molecules of the sample must undergo dipole moment change during
vibration. When the frequency of the IR is the same as the vibrational frequency of
the bonds, absorption occurs and a spectrum can be recorded. Different functional
groups absorb heat at different frequencies dependent upon their structure, and
thus a vibrational spectrum can be used to determine the functional groups present
in a sample. When interpreting the data obtained by an IR, results can be compared
to a frequency table to find out which functional groups are present to help
determine the structure.

3. Raman Spectroscopy: Raman spectroscopy is similar to IR in that it is a vibrational
spectroscopy technique, but it uses inelastic scattering. The spectrum of Raman

spectroscopy shows a scattered Rayleigh line and the Stoke and anti-Stoke lines,
which is different to the irregular absorbance lines of IR. Raman spectroscopy works
by the detection of inelastic scattering, also known as Raman scattering, of
monochromatic light from a laser in the visible, near infrared or ultraviolet range. For
a transition to be Raman active, there must be a change in the polarizability of the
molecule during the vibration and the electron cloud must experience a positional
change. The technique provides a molecular fingerprint of the chemical composition
and structures of samples, but Raman scattering gives inherently weak signals.
Techniques such as Surface Enhanced Raman Spectroscopy (SERS) have been
developed to enhance sensitivity when using Raman spectroscopy.

4. Nuclear Magnetic Resonance (NMR): Nuclear magnetic resonance (NMR) uses
resonance spectroscopy and nuclear spin states for spectroscopic analysis. All nuclei
have a nuclear spin, and the spin behaviour of the nucleus of every atom depends on
its intramolecular environment and the external applied field. When nuclei of a
particular element are in different chemical environments within the same molecule,
there will be varied magnetic field strengths experienced due to shielding and de-
shielding of electrons close by, causing different resonant frequencies and defines
the chemical shift values. Spin-spin coupling takes into account that the spin states
of one nucleus affects the magnetic field that is experienced by neighbouring nuclei,
via intervening bonds. Spin-spin coupling causes absorption peaks of each group of
nuclei to be split into a number of components. There are multiple types of NMR
analysis, which are hydrogen NMR, carbon 13 NMR, DEPT 90 and DEPT 135 NMR.
The NMR spectrum of a compound shows the resonance signals that are emitted by
the atomic nuclei present in a sample, and these can be used to identify the
structure of a compound.

1.4.3 Selection Rules for Electronic Transition

1. The total spin cannot change, ΔS=0; the rule ΔΣ=0 holds for multiplets; If the spin-
orbit coupling is not large, the electronic spin wavefunction can be separated from
the electronic wavefunctions. Since the electron spin is a magnetic effect, electronic
dipole transitions will not alter the electron spin. As a result, the spin multiplicity
should not change during the electronic dipole transition.

2. The total orbital angular momentum change should be ΔA=0, ±1; For heteronuclear
diatomic molecules with C∞v symmetry, a Σ +↔ Π transition is allowed according to
ΔA selection rule. In order to prove the allowance of this transition, the direct
product of Σ+ ⊗Π yields Π irreducible representation from the direct product table.
Based on the C∞v character table below, the operator in the x and y direction have

doubly degenerate Π symmetry. Therefore, the transition between Σ +↔ Π must be
allowed since the multiplication of any irreducible representation with itself will
provide the totally symmetric representation. The electronic transition moment
integral can be nonzero. We can use the same kind of argument to illustrate that
\(\Sigma ^ - \leftrightarrow \Phi\) transition is forbidden.

3. Parity conditions are related to the symmetry of the molecular wavefunction
reflecting against its symmetry axis. For homonuclear molecules, the g ↔ u
transition is allowed. For heteronuclear molecules, + ↔ + and - ↔ - transitions
apply; For hetero diatomic molecules with C∞v symmetry, we can use group theory
to reveal that Σ+ ↔Σ+ and Σ− ↔Σ−transitions are allowed, while Σ+ ↔Σ−
transitions are forbidden. The direct product of either Σ+ ⊗Σ+ or Σ− ⊗Σ− yields the
same irreducible representation Σ + based on the direct product table. The z
component of the dipole operator has Σ + symmetry. The transition is allowed
because the electronic transition moment integral can generate the totally
symmetric irreducible representation Σ + . We can use the same method to prove
that Σ+ ↔Σ− transitions are forbidden. Similarly, for a molecule with an inversion
centre, a subscript g or u is used to reveal the molecular symmetry with respect to
the inversion operation, i. The x, y, and z components of the transition dipole
moment operator have u inversion symmetry in molecule with inversion centre (g
and u are short for gerade and ungerade in German, meaning even and odd). The
electronic transition moment integral need to yield totally symmetric irreducible
representation Ag for allowed transitions. Therefore, only g ↔ u transition is
allowed.

1.4.4 Types of Molecular Energies

1. Translational energy levels: The
translational energy levels of a
molecule are usually taken to be
those of a particle in a three-
dimensional box:

In general, the separation of the translational energy levels is many orders of
magnitude smaller than kT, even for a box with dimensions on the order of
molecular sizes. The translational behaviour of molecules therefore appears to be
classical.
2. Rotational energy levels – diatomic molecules: Diatomic molecules are often
approximated as rigid rotors, meaning that the bond length is assumed to be fixed.
Solving the Schrodinger equation for a rigid rotor gives the following energy levels:

In this equation, J is the quantum number for total rotational angular momentum,
and B is the rotational constant, which is related to the moment of inertia , I = r 2 (
is the reduced mass and r the bond length) of the molecule.

B is usually given in units of cm1, and in order to end up with these units, the speed
of light in this equation, c, needs to be in cm s1 rather than ms1. Using B in cm1 will
yield the energy in the same units. [Note: cm1 are common units in spectroscopy for
historical reasons, even though they are non-SI. Units of m1 are never used, so don’t
be tempted to convert, however stupid it might seem...] The equation above for E(J)
is only approximate. In order to provide a better description of the energy levels of a
diatomic molecule, a centrifugal distortion term is added to the energy. This takes
account of the fact that as a real molecule rotates faster and faster (i.e. with more
energy), the bond stretches a little (if you swing a weight attached to a piece of
elastic in a circle, you can observe the same effect on a macroscopic scale).

D is called the centrifugal distortion constant, and is several orders of magnitude
smaller than B.

Separation of rotational energy levels correspond to the microwave region of the
electromagnetic spectrum. In order for a molecule to absorb a microwave radiation,
it must have a permanent dipole moment, i.e., µ≠0. This condition is known as the
gross selection rule for microwave, or pure rotational, spectroscopy. Specific
selection rules arise largely from conservation of angular momentum, and generally
involve statements of allowed changes in quantum number. In general, the selection
rule for changes in rotational angular momentum, following absorption of photon is
J=0, ±1. However, for the rotational transitions, J=0 does not correspond to a
transition since the particle stays in the same state, so the specific selection rule for a
pure rotation spectrum is J=±1.

The transition energies may be calculated from the energy equation above. For a
transition from level J to level J+1, we have:

Neglecting the centrifugal distortion term, the spacing between the lines in a
spectrum will be

i.e. lines in a pure rotational spectrum are (almost) equally spaced by 2B (almost
because we have neglected the effects of centrifugal distortion). This result is often
useful if we are attempting to assign lines in a spectrum in terms of the quantum
numbers of the states involved. The separation between a pair of lines allows us to
estimate B (i.e. B=separation/2), and we can use E(J) = 2B(J+1) to determine the J
values involved in a given transition.

3. Rotational energy levels - polyatomic molecules: Polyatomic molecules may
rotate about the x, y or z axes, or some combination of the three. They have
moments of inertia Ix, Iy, Iz associated with each axis, and also corresponding
rotational constants A, B and C
[A = h/(82cIx), B = h/(82cIy), C = h/(82cIz)]. The general equation for a moment of
inertia is:

where the sum is over the atoms in the molecule, mi is the mass of atom i and ri is its
distance from the rotation axis. [Note: this explains why only two rotation axes are
counted in a diatomic molecule. Rotation about the principal axis would have I=0,
since all the atoms lie along the axis, and since the kinetic energy associated with
rotation is K = ½ I2 , with  the angular frequency of rotation, there would be no
energy associated with such a rotation (and therefore no energy levels and no
transitions between energy levels).] In general, the rotational constants A, B, and C
may all be different, and a molecule for which this is true is called an asymmetric top.
The spectroscopy of such molecules is quite complicated, and beyond the scope of
the course. Instead, we will look at symmetric top molecules, such as CH3F, CH3Cl,
CCl3H or benzene, for which two of the rotational constants are equal. Symmetric
tops may be divided into prolate tops, for which B > A, C (e.g. CH3F, CH3Cl) and oblate
tops, for which B < A, C (e.g. CCl3H, benzene). The rotational motion of a symmetric
top molecule is described by three quantum numbers:

J – the total angular momentum
K – the projection of J onto the principal axis of the molecule
MJ – the projection of J onto some chosen lab-frame z axis (often just written M)

Regardless of which category the molecule falls into, the rotational energy is given
approximately by:

When centrifugal distortion is included, this becomes

The selection rules for a symmetric top molecule are J = ±1, K = 0.
Transition energies and line separations may be calculated in a similar way to that
worked through above for diatomic molecules, by determining E(J+1, K) – E(J,K), etc.
3. Vibrational energy levels
To a first approximation, molecular vibrations can be approximated as simple
harmonic oscillators, with an associated energy

where v is the vibrational quantum number and  is the vibrational frequency (the
symbols look quite similar – don't confuse them!). In the above expression, E(v) is in
units of Joules. Usually, the energies are given in cm1, in which case h is replaced
by e, the vibrational wavenumber (i.e. this is simply the harmonic vibrational
frequency or energy expressed in units of cm1).

This is usually a fairly good approximation near the bottom of the potential well,
where the potential closely resembles that of a harmonic oscillator. However, real
molecules have enharmonic potentials, and this is taken into account by rewriting
the energy as a series expansion, with successively higher order correction terms:

Often, only the first correction term, involving xee, will be used, but note that the
source of differences in calculated and measured properties involving vibrational
energy levels is frequently neglect of higher order terms. Note that every vibration
has a zero point energy found by substituting v=0 into the above equation, ZPE = E(0)
= ½e – ¼xee.
As well as changing the energy of the vibrational levels, anharmonicity has another,
less obvious effect: for a molecule with an anharmonic potential, the rotational
constant changes slightly with vibrational state. The rotational constant Bv for a given
vibrational state can be described by the expression:

where Be is the rotational constant corresponding to the equilibrium geometry of the
molecule, e is a constant determined by the shape of the anharmonic potential, and
v is the vibrational quantum number.
Returning now to considering the vibrational energy levels of a harmonic and an
anharmonic oscillator, recall that the energy levels in a harmonic oscillator are

equally spaced by an energy e. The effect of anharmonicity is that the levels get
increasingly closer together as the vibrational quantum number v increases, until
eventually there is no difference in energy between successive quantum states, and
the energy levels form a continuum. The energy at which this occurs (known as the
dissociation energy) coincides with the point at which the potential energy function
reaches zero and the molecule is no longer bound (this is the reason why the energy
levels are no longer quantized). The dissociation energy may be measured either
from the bottom of the potential, in which case it is called De, or from the zero-point
energy, in which case it is called D0.
To determine the dissociation energy, we set dE/dv equal to zero to determine the
maximum value of the vibrational quantum number v.

Substituting back into the energy expression above gives

De and D0 are related by

Let us Revise:

1. There are how many types of Molecular Energies?
2. What is NMR Spectroscopy?
3. Explain Spin Rule for Electronic Transition?
4. How is Raman spectroscopy similar to IR spectroscopy?
5. Explain UV and Visible Spectroscopy?

Summary

Spectroscopy: Large number of techniques using electromagnetic radiation to obtain
information of structure and nature of matter.
Electromagnetic Radiations: Generated by acceleration of charged particles.
Electromagnetic Spectrum: band of enormous range of wavelengths and frequencies
of electromagnetic radiations generated due acceleration of charged particles.
Generally divided into seven regions.

Classification of Spectroscopy

|

_________|_________

||

Atomic Molecular

_______________________|______ ________|_____________________

|| | || | | |

Emission Absorption | Florescent UV, Raman Infrared NMR

Phosphorescent Visible

3 Selection Rules for Electronic Transition: Spin, Angular Momentum, Symmetry.
4 types of molecular energies depending on Degrees of Freedom:
1. Translational
2. Rotational-Diatomic
3. Rotational-Triatomic
4. Vibrational

Short Answer Questions

1. Write 3 points of comparison between UV and IR Spectrum.
2. Explain the characteristics of Electromagnetic Radiations.
3. Explain the principles of Spectroscopy.
4. List down various constituents of the Electromagnetic Spectrum.

5. What is the purpose of an IR Spectrometer?
6. Can X-rays be used instead of microwaves to study pure rotational spectra?
7. What property of nuclei is involved in NMR spectroscopy?
8. What type of excitation can take place in CH3NO2 in UV-spectroscopy?
9. The presence of hydrogen atoms attached to an aromatic ring may be identified by

UV spectrometer. State true or false. Give reason.
10. Name the spin active nuclei which can be studied by NMR spectroscopy.

Long Answer Questions

1. Explain each constituent of the Electromagnetic Spectrum.
2. Explain various terms to energy changes in a molecule.
3. What are the Selection Rules for Electronic Transitions? Explain briefly.
4. Explain the different types of Molecular Spectroscopy.
5. Discuss each type of Molecular Energies.
6. IR spectra is often characterised as “molecular finger prints”, Justify this statement.
7. What type of excitation can take place in CH3NO2 in UV-spectroscopy?
8. What are limitations of microwave spectra?
9. Give a schematic diagram for typical IR spectrometer.
10. What is the purpose of NMR spectroscopy?

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

2

Application of Spectroscopy

2.1 Introduction

Spectroscopy also provides a precise analytical method for finding the constituents in material
having unknown chemical composition. In a typical spectroscopic analysis, a concentration of a
few parts per million of a trace element in a material can be detected through its
emission spectrum.

In astronomy the study of the spectral emission lines of distant galaxies led to the discovery that
the universe is expanding rapidly and isotopically (independent of direction). The finding was
based on the observation of a Doppler shift of spectral lines. The Doppler shift is an effect that
occurs when a source of radiation such as a star moves relative to an observer.
The frequency will be shifted in much the same way that an observer on a moving train hears a
shift in the frequency of the pitch of a ringing bell at a railroad crossing. The pitch of the bell
sounds higher if the train is approaching the crossing and lower if it is moving away. Similarly,
light frequencies will be Doppler-shifted up or down depending on whether the light source is
approaching or receding from the observer. During the 1920s, the American astronomer Edwin
Hubble identified the diffuse elliptical and spiral objects that had been observed as galaxies. He
went on to discover and measure a roughly linear relationship between the distance of these
galaxies from Earth and their Doppler shift. In any direction one looks, the farther the galaxy
appears, the faster it is receding from Earth.

Spectroscopic evidence that the universe was expanding was followed by the discovery in 1965
of a low level of isotropic microwave radiation by the American scientists Arno A.
Penzias and Robert W. Wilson. The measured spectrum is identical to the radiation
distribution expected from a blackbody, a surface that can absorb all the radiation incident on it.
This radiation, which is currently at a temperature of 2.73 kelvin (K), is identified as a relic of
the big bang that marks the birth of the universe and the beginning of its rapid expansion.

2.2 Emission Spectroscopy

Classical emission spectroscopy is based on excitation of atoms or molecules into higher
electronic states by electron impact (in gas discharges), photon absorption or thermal
excitation at high temperatures (in star atmospheres). Excitation by narrow-band lasers may
result in the selective population of wanted levels, which emit their excitation energy as
fluorescence photons.

Such a laser-induced fluorescence (LIF) spectrum is much simpler than the emission spectrum of
a gas discharge, where the superposition of fluorescence from many emitting levels is observed.

Measurements of spectrally resolved LIF allows the determination of the final levels into which
the fluorescence is emitted. This gives access, for example, to high vibrational–rotational levels
in the electronic ground state that are not thermally populated and are therefore not accessible

to absorption spectroscopy. Examples are levels closely below the dissociation energy, where
the vibrating atoms reach very large internuclear distances at their outer turning point.

As in AA spectroscopy, the sample must be converted to free atoms, usually in a high-
temperature excitation source. Liquid samples are nebulized and carried into the excitation
source by a flowing gas. Solid samples can be introduced into the source by a slurry or by laser
ablation of the solid sample in a gas stream. Solids can also be directly vaporized and excited by
a spark between electrodes or by a laser pulse. The excitation source must desolvate, atomize,
and excite the analyte atoms. A variety of excitation sources are described in separate
documents:

Direct-current plasma (DCP) Flame

Inductively-coupled plasma (ICP) Laser-induced breakdown (LIBS) Laser-induced plasma
Microwave-induced plasma (MIP) Spark or arc

Direct-Current Plasma Excitation Source

A direct-current plasma (DCP) is created by an electrical discharge between two electrodes. A
plasma support gas is necessary, and Ar is common.

Samples can be deposited on one of the electrodes, or if conducting can make up one electrode.
Insulating solid samples are placed near the discharge so that ionized gas atoms sputter the
sample into the gas phase where the analyte atoms are excited. This sputtering process is often
referred to as glow-discharge excitation.

Inductively-Coupled Plasma (ICP) Excitation Source

An inductively coupled plasma (ICP) is a very high temperature (7000-8000K) excitation source
that efficiently desolvates, vaporizes, excites, and ionizes atoms. Molecular interferences are
greatly reduced with this excitation source but are not eliminated completely. ICP sources are
used to excite atoms for atomic-emission spectroscopy and to ionize atoms for mass
spectrometry.

Spark and Arc Emission Sources

Spark and arc excitation sources use a current pulse (spark) or a continuous electrical
discharge (arc) between two electrodes to vaporize and excite analyte atoms. The electrodes
are either metal or graphite. If the sample to be analysed is a metal, it can be used as one
electrode. Non-conducting samples are ground with graphite powder and placed into a cup-
shaped lower electrode. Arc and spark sources can be used to excite atoms for atomic- emission
spectroscopy or to ionize atoms for mass spectrometry. Arc and spark excitation sources have
been replaced in many applications with plasma or laser sources, but are still widely used in the
metals industry.

2.2.1 Flame Photometry

Principle: The compounds of the alkali and alkaline earth metals (Group II) dissociate into
atoms when introduced into the flame. Some of these atoms further get excited to even higher
levels. But these atoms are not stable at higher levels.

Hence, these atoms emit radiations when returning back to the ground state. These radiations
generally lie in the visible region of the spectrum. Each of the alkali and alkaline earth metals
has a specific wavelength.

For certain concentration ranges,
The intensity of the emission is directly proportional to the number of atoms returning to the
ground state. And the light emitted is in turn proportional to the concentration of the sample.

If a solution containing a metallic salt is aspirated into a flame, a vapour which contains atoms of
the metal may be formed. Some of these gaseous metal atoms may be raised to an energy level
which is sufficiently high to permit the emission of radiation characteristic of the metal, e.g., its
characteristic yellow colour imparted to flames by compounds of sodium. This is the basic of
Flame Emission Spectroscopy (FES) formerly known as Flame Photometry.
The procedure by which gaseous metal atoms are produced in the flame may be summarized as
follows:
When a solution containing a suitable compound of the metal to be investigated is aspirated
into a flame, the following events occur in rapid succession:

1. Evaporation of solvent, leaving a solid residue.
2. Vaporization of the solid with dissociation into its constituent atoms, initially in the

ground state.
Some atoms may be excited by the thermal energy of the flame to higher energy levels and
attain a condition in which they radiate energy.

2.2.1.1 Flame Photometer

The flame is a much lower energy source than the electrical means of excitation. Thus, the
flame produces a simpler emission spectrum with fewer lines. Furthermore, relatively cool
frames are normally used in flame photometers. In a flame photometer the emitted
radiation is isolated by an optical fiber and then converted and then converted to an
electrical signal by the photodetector.

Basic Components: Flame emission spectroscopy requires a nebulizer-burner system, which
produces gaseous metal atoms by using a suitable combustion flame involving a fuel gas-
oxidant gas mixture.
Flame temperature: An essential requirement of flame spectroscopy is a flame temperature
greater than 2000K. in most cases, this can only be met by burning a fuel gas in an oxidant
gas, which is usually air, nitrous oxide or oxygen diluted with either nitrogen or argon. An
acetylene-air mixture is suitable for the determination of some 30 metals, but a propane-air
flame is preferred for metals, which are easily converted into an atomic vapour state.
Method: Air at a given temperature is passed into an atomiser and the suction, thus
produced, draws a solution of the sample into the atomiser, where it joins the air steam as a
fine mist and passes into the burner at a given pressure and the mixture is burnt. Radiation
form the resulting flame passes through a lens and finally through an optical fibre which

permits only the radiation
characteristic of the element
under investigation to pass
through the photodetector.

Evaluation Methods: In order
to convert the measured
emission values into the
concentration of the analyte,
the following methods may be
used:

a. Calibration curve.
b. The internal standard method involves the addition of a fixed amount of reference

material to the sample solution and the standard solution. Upon excitation, the
emitted energy of the analyte element and the internal standard are measured
simultaneously by two photodetectors.

Applications of flame photometer

1. Flame photometer can be applied both for quantitative and qualitative analysis of
elements. The radiations emitted by the flame photometer are characteristic to
particular metal. Hence with the help of Flame photometer we can detect the
presence of any specific element in the given sample.

2. The presence of some group II elements is critical for soil health. We can determine
the presence of various alkali and alkaline earth metals in soil sample by conducting
flame test and then the soil can be supplied with specific fertiliser.

3. The concentrations of Na+ and K+ ions are very important in the human body for
conducting various metabolic functions. Their concentrations can be determined by
diluting and aspirating blood serum sample into the flame.

4. Soft drinks, fruit juices and alcoholic beverages can also be analysed by using flame
photometry to determine the concentrations of various metals and elements.

Advantages:

1. Simple quantitative analytical test based on the flame analysis.
2. Inexpensive.
3. The determination of elements such as alkali and alkaline earth metals is performed

easily with most reliable and convenient methods.
4. Quite quick, convenient, and selective and sensitive to even parts per million (ppm)

to parts per billion (ppb) range.

Disadvantages:

Moreover, the flame photometer has a wide range of applications in the analytical
chemistry, it possesses many disadvantages which are explained below:

1. The concentration of the metal ion in the solution cannot be measured accurately.
2. A standard solution with known molarities is required for determining the

concentration of the ions which will corresponds to the emission spectra.
3. It is difficult to obtain the accurate results of ions with higher concentration.
4. The information about the molecular structure of the compound present in the

sample solution cannot be determined.
The elements such as carbon, hydrogen and halides cannot be detected due to its non-
radiating nature.

2.3 Fluorescence Spectroscopy

Fluorescence spectroscopy analyses fluorescence from a molecule based on its fluorescent
properties.

Fluorescence is a type of luminescence caused by photons exciting a molecule, raising it to
an electronic excited state.
Fluorescence spectroscopy uses a beam of light that excites the electrons in molecules of
certain compounds, and causes them to emit light. That light is directed towards a filter and
onto a detector for measurement and identification of the molecule or changes in the
molecule.

Steady state fluorescence spectra are when molecules, excited by a constant source of light,
emit fluorescence, and the emitted photons, or intensity, are detected as a function of
wavelength. A fluorescence emission spectrum is when the excitation wavelength is fixed
and the emission wavelength is scanned to get a plot of intensity vs. emission wavelength.
A fluorescence excitation spectrum is when the emission wavelength is fixed and the
excitation monochromator wavelength is scanned. In this way, the spectrum gives
information about the wavelengths at which a sample will absorb so as to emit at the single
emission wavelength chosen for observation. It is analogous to absorbance spectrum, but is
a much more sensitive technique in terms of limits of detection and molecular specificity.
Excitation spectra are specific to a single emitting wavelength/species as opposed to an
absorbance spectrum, which measures all absorbing species in a solution or sample. The
emission and excitation spectra for a given fluorophore are mirror images of each other.
Typically, the emission spectrum occurs at higher wavelengths (lower energy) than the
excitation or absorbance spectrum.
Emission after excitation with UV light

These two spectral types (emission and excitation) are used to see how a sample is
changing. The spectral intensity and or peak wavelength may change with variants such as
temperature, concentration, or interactions with other molecules around it. This includes
quencher molecules and molecules
or materials that involve energy
transfer. Some fluorophores are also
sensitive to solvent environment
properties such as pH, polarity, and
certain ion concentrations.
Fluorescent molecules and materials
come in all shapes and sizes. Some
are intrinsically fluorescent, such as
chlorophyll and the amino acid
residue tryptophan (Trp),
phenylalanine (Phe) and tyrosine
(Tyr). Others are molecules

synthesized specifically as stable organic dyes or tags to be added to otherwise non-
fluorescent systems. There are entire catalogues of these available. Typically, organic
fluorescent molecules have aromatic rings and pi- conjugated electrons in them. Depending
on their size and structure, organic dyes can emit from the UV out into the near-IR.

Here is a random sampling of a few common fluorophores that span the UV and Visible
range. Some rare earth elements, or lanthanides, have higher electronic orbitals filled,
where electrons transition due to metal ligand charge transfers happen between 4f-5d and
even 4f-4f orbitals. (Bunzli, 1989) There are many molecules that are luminescent in nature
such as a few of the amino acids, chlorophylls, and natural pigments. Others are highly
engineered for very specific uses of fluorescence spectroscopy.

A few of the categories of fluorescent molecules and materials are:
Amino acids (Trp, Phe, Tyr)

Base pair derivatives (2-AP, 3-MI, 6-MI, 6-MAP, pyrrolo-C, tC) Chlorophylls

Fluorescent Proteins (FPs)
Organic dyes
(fluorescein,
rhodamine, N-
aminocoumarins and
derivatives of these)
Rare earth elements
(lanthanides)
Semiconductors
Quantum dots
Single Walled Carbon
Nanotubes (SWCNTs)
Solar cells
Pigments, brighteners
Phosphors
Many more…

Other molecules and
materials such as
fluorescent proteins,
semiconductors,
phosphors, and rare
earth elements are
among the commonly
used fluorescent
samples. Polymers
with conjugated aromatics or dienes also commonly have fluorescent properties. Of course,
new materials are being created all the time.

2.4 Phosphorescence Spectroscopy

2.4.1 Fluorescence and phosphorescence

These phenomena are closely related to electronic absorption spectra and can be used as a
tool for analysis and structure determination. Both involve the absorption of radiation via an
electronic transition, a loss of energy through either vibrational energy decay or
nonradiative processes, and the subsequent emission of radiation of a lower frequency than
that absorbed. Electrons possess intrinsic magnetic moments that are related to their spin
angular momenta. The spin quantum number is s = 1/2, so in the presence of a magnetic
field an electron can have one of two orientations corresponding to magnetic spin quantum
number ms = ±1/2. The Pauli exclusion principle requires that no two electrons in an atom
have the same identical set of quantum numbers; hence when two electrons reside in a
single AO or MO they must have different ms values (i.e., they are antiparallel, or spin
paired).This results in a cancellation of their magnetic moments, producing a so-called
singlet state. Nearly all molecules that contain an even number of electrons have singlet
ground states and have no net magnetic moment (such species are called diamagnetic).
When an electron absorbs energy and is excited to a higher energy level, there exists the
possibility of (1) retaining its antiparallel configuration relative to the other electron in the
orbital from which it was promoted so that the molecule retains its singlet characteristic, or
(2) changing to a configuration in which its magnetic moment is parallel to that of its
original paired electron.

In the latter case, the molecule will possess a net magnetic moment (becoming
paramagnetic) and is said to be in a triplet state. For each excited electronic state, either
electron spin configuration is possible so that there will be two sets of energy levels. The
normal selection rules forbid transitions between singlet (Si) and triplet (Ti) states; hence
there will be two sets of electronic transitions, each associated with one of the two sets of
energy levels.

2.4.2 Fluorescence

Fluorescence is the process whereby a molecule in the lower of two electronic states
(generally the ground state) is excited to a higher electronic state by radiation whose

energy corresponds to an allowed absorption transition, followed by the emission of
radiation as the system decays back to the original state. The decay process can follow
several pathways. If the decay is back to the original lower state, the process is called
resonance fluorescence and occurs rapidly, in about one nanosecond. Resonance
fluorescence is generally observed for monatomic gases and for many organic molecules, in
particular aromatic systems that absorb in the visible and near-ultraviolet regions. For many
molecules, especially aromatic compounds whose electronic absorption spectra lie
predominately in the shorter-wavelength ultraviolet region (below 400 nanometres), the
lifetime of the excited electronic state is sufficiently long that prior to the emission of
radiation the molecule can (1) undergo a series of vibrational state decays, (2) lose energy
through interstate transfer (intersystem crossing), or
(3) lose vibrational energy via molecular collisions.
In the first case, the system will emit radiation in the infrared region as the vibrational
energy of the excited state decays back to the lowest vibrational level. The molecule then
undergoes an electronic state decay back to one of the vibrational states associated with
the lower electronic state.

The resulting emission spectrum will then be centred at a frequency lower than the
absorption frequency and will appear to be a near mirror image of the absorption spectrum.
The second mechanism can be illustrated by reference to the potential energy curves for
nitrogen hydride (NH) shown in Figure 7B. The curves for the 1Σ+ and 1Π states intersect at
a radius value of 0.2 nanometre. If a molecule in the 1Π excited electronic state is in a
vibrational level corresponding to the energy value of this intersection point, it can cross
over to the 1Σ+ state without emission or absorption of radiation. Subsequently it can
undergo vibrational energy loss to end up in the lowest vibrational state of the 1Σ+
electronic state. This can then be followed by an electronic transition back to the lower 1Δ
state. Thus, the absorption of energy corresponding to an original 1Δ → 1Π transition results
in the emission of fluorescence radiation corresponding to the lower frequency 1Σ+ → 1Δ
transition. In the third case, when two molecules collide there exists the possibility for
energy transfer between them. Upon colliding, a molecule can thus be transformed into a

different electronic state whose energy minimum may lie lower or higher than its previous
electronic state.

The lifetimes of the excited singlet electronic states, although long enough to allow
vibrational relaxation or intersystem crossing, are quite short, so that fluorescence occurs
on a time scale of milliseconds to microseconds following irradiation of a material. The most
common mode of observation of fluorescence is that of using ultraviolet radiation (invisible
to the human eye) as an exciting source and observing the emission of visible radiation. In
addition to its use as a tool for analysis and structural determination of molecules, there are
many applications outside the laboratory. For example, postage stamps may be tagged with
a visually transparent coating of a fluorescing agent to prevent counterfeiting, and the
addition of a fluorescing agent with emissions in the blue region of the spectrum to
detergents will impart to cloth a whiter appearance in the sunlight.

2.4.3 Phosphorescence

Phosphorescence is related to fluorescence in terms of its general mechanism but involves
a slower decay. It occurs when a molecule whose normal ground state is a singlet is
excited to a higher singlet state, goes to a vibrationally excited triplet state via either an
intersystem crossing or a molecular collision, and subsequently, following vibrational
relaxation, decays back to the singlet ground state by means of a forbidden transition. The
result is the occurrence of a long lifetime for the excited triplet state; several seconds up to
several hours are not uncommon.

These long lifetimes can be related to interactions between the intrinsic (spin) magnetic
moments of the electrons and magnetic moments resulting from the orbital motion of the
electrons.

Molecules in singlet and triplet states react chemically in different manners. It is possible to
affect chemical reactions by the transfer of electronic energy from one molecule to another
in the reacting system. Thus, the study of fluorescence and phosphorescence provides
information related to chemical reactivity.

Phosphorescence is luminescence that is emitted for a significant period of time (in the
millisecond to second time regime) after excitation.

2.4.4 Applications of Florescence Spectroscopy

Laser induced fluorescence spectroscopy of human
tissues for cancer diagnosis: Cancer is one of the most
dreaded diseases of our time and has a very high
incidence. Early tumours often arise from tissue which
have a rapid turnover of cells and are active in repair like
transformed mucosa on the surface of hollow organs
(oral cavity, gastrointestinal tract, female reproductive
organs etc.). Laser spectroscopic techniques have the
potential for in-situ, near real time diagnosis and the use

of non-ionizing radiation ensures that the diagnosis can be made repeatedly without any
adverse side effects.

Laser Induced Fluorescence (LIF) has been used for diagnosing cancer in two ways. One
approach involves systemic administration of a drug like hematoporphyrin derivative (HpD)
which is selectively retained by the tumour. When photo excited with light of appropriate
wavelength the drug localized in the tumour fluoresces. This fluorescence is used for
detection and imaging of the tumour. Photo excitation also leads to populating the triplet
state via intersystem crossing. The molecule in excited triplet state can directly react with
bio-molecules or lead to generation of singlet oxygen which is toxic to the host tissue. The
resulting destruction of the host tissue is exploited for photodynamic therapy of tumour.

Study of Marine Petroleum Pollutants: Fluorescence spectroscopy is one of the good
techniques to detection of oil slicks on the water surface, determination of petroleum
contaminants in seawater and determination of particular petroleum derivative compounds
as well as identification of pollution sources. Main components of any oil are hydrocarbons.
The other components are primarily derivatives of hydrocarbons containing single atoms of
sulfur, oxygen or nitrogen. Only a few of hydrocarbons fluoresce, while the major of them
show no ability to luminescence. The content of compounds able to fluorescence rarely
exceeds 10% of the oil mass. At the same time the petroleum strongly absorbs radiation,
especially the ultraviolet and blue light. In spite of this petroleum is a luminescent medium
and fluorescence is a phenomenon which allows testing oils. Fluorescence of oils has
wavelength over then 260 nm and covers a spectral area of ultraviolet and visible light. The
phenomenon is most significant in the 270–400 nm range.

Accurate determination of glucose: Glucose is considered as a major component of animal
and plant carbohydrates in biological systems. Furthermore, blood glucose levels are also an
indicator of human health conditions: the abnormal amount of glucose provides significant
information of many diseases such as diabetes or hypo glycemia. Florophotometry was used
widely owing to its operational simplicity and high sensitivity. Recently bio-molecule-
stabilized Au nanoclusters were demonstrated as a novel fluorescence probe for sensitive
and selective detection of glucose.

Future Research

Spectroscopic technique may be automatized which can then process many diagnostic
samples at the same time. Also, fibre optic systems may be integrated with this
spectroscopic technique to diagnose microorganisms in vivo. By this modification, infections
in many body parts can be detected with ease. Further research is required to establish
flexible and portable spectroscopic devices which can be integrated in daily medical
practice.

There is need for reference libraries for spectral signatures of individual microorganism. This
will be very helpful for comparison with spectral signatures from an unknown
microorganism sample. But there are many questions which remain to be answered like if
biological sample contains more than one microorganism, then how it will affect the
spectral signature appearance and how to interpret these spectral for making definite

diagnosis. Also, microorganisms like bacteria have many chemicals which are same like in
human cells and in extracellular space, thus body fluids samples may contain same
chemicals as found in microorganisms. As a result, it may interfere with spectroscopic
spectral analysis and may be a hurdle to reach on definite diagnosis. This justifies the need
for studies which can enable to make distinction between microorganism and human cells.
Also, future studies should be directed to determine the specific spectral regions which will
be suitable for identification of specific microorganisms. It will help to design invasive and
non-invasive techniques for microorganism's diagnosis inside the body cavities by use of
fibre optic devices.

2.5 Jablonski Diagram

In molecular spectroscopy, a Jablonski diagram is
a diagram that illustrates the electronic states of
a molecule and the transitions between them.
The states are arranged vertically by energy and
grouped horizontally by spin multiplicity.

Nonradiative transitions arise through several
different mechanisms, all differently labelled in
the diagram. Relaxation of the excited state to its
lowest vibrational level is called vibrational
relaxation. This process involves the dissipation of
energy from the molecule to its surroundings,
and thus it cannot occur for isolated molecules. A
second type of nonradiative transition is Internal
Conversion (IC), which occurs when a vibrational
state of an electronically excited state can couple
to a vibrational state of a lower electronic state.
A third type is InterSystem Crossing (ISC); this is a
transition to a state with a different spin
multiplicity. In molecules with large spin-orbit
coupling, intersystem crossing is much more important than in molecules that exhibit only
small spin-orbit coupling. ISC can be followed by Phosphorescence.

2.5.1 Radiative and Non-Radiative Transitions

The radiative and non-radiative transitions that lead to the observation of
molecular photoluminescence are typically illustrated by an energy level diagram called the
Jablonski diagram. Figure shows a Jablonski diagram that explains the mechanism of light
emission in most organic and inorganic luminophores. The spin multiplicity of a given
electronic state can be either a singlet (paired electrons) or a triplet (unpaired electrons).
The ground electronic state is normally a singlet state and is designated as S0 in Figure Excited
electronic states are either singlet (S1, S2) or triplet (T1) states.

The Jablonski diagrams. Radiative and non-radiative processes are depicted as solid and
dashed lines, respectively.

When the molecule absorbs light an electron is promoted within 10−14–10−15 s from the
ground electronic state to an excited state that should possess the same spin multiplicity as
the ground state. This excludes a triplet excited state as the final state of electronic
absorption because the selection rules for electronic transitions dictate that the spin state
should be maintained upon excitation. A plethora of non-radiative and radiative processes
usually occur following the absorption of light enroot to the observation of
molecular luminescence.

2.5.1.1 How to Interpret a Jablonski Diagram

The Jablonski diagram is a powerful tool for visualising the possible transitions that can
occur after a molecule has been photoexcited. A typical Jablonski diagram is shown and
the key components and transitions that make up the diagram are explained below.

Energy Levels
The energy levels of a molecule are shown by the horizontal black lines; with energy
increasing along the vertical axis of the diagram. The bold lines represent the lowest
vibrational level of each electronic state, with the higher vibrational levels represented by
thinner lines. The vibrational levels become more closely spaced as energy increases and
eventually form a continuum; for clarity, only a subset of these vibrational levels are
represented on the diagram.

The naming of the electronic states is based on the spin angular momentum configuration
of each state. Singlet states (a total spin angular momentum of zero) are denoted by
an S and triplet states (a total spin angular momentum of one) by T:

S0 is the singlet ground state of the molecule
S1 is the first excited singlet state and Sn is the nth excited singlet state
T1 is the first excited triplet state and Tn is the nth excited triplet state
Radiative & Non-Radiative Transitions
The coloured arrows represent the various transitions that can transfer energy between
the molecular states and are split into radiative and non-radiative transitions.

Radiative transitions are transitions between two molecular states where the energy
difference is emitted or absorbed by photons and are represented in a Jablonski diagram
by straight arrows.

Non-radiative transitions are transitions between two molecular states without the
absorption or emission of photons and are represented in a Jablonski diagram by
undulating arrows.

Summary

Emission Spectroscopy: based on excitation of atoms or molecules from ground to
excited state which when fall down emit energy equal to excitation energy.
Flame Photometry:
Principle: The compounds of the alkali and alkaline earth metals (Group II) dissociate
into atoms when introduced into the flame, acquire energy to higher levels, come
back due to instability to release radiations in the Visible Region of EM Spectrum.
Florescence spectroscopy: When excited atoms or molecules come directly to lower
energy levels without going to Tn, without Intersystem Crossing (ISC), this results in
florescence.
Phosphorescence: when excited molecules undergo Intersystem Crossing (ISC) from
Sn to Tn and then to lower energy levels.
Various applications of Florescent Spectroscopy in the fields of medical sciences,
engineering, agriculture, etc.
Jablonski Diagram: powerful tool for visualising the possible transitions that can
occur after a molecule has been photoexcited.
Energy States: Singlet States (Sn) and Triplet States (Tn).