Introduction
The
molecular spectroscopy is the study of the interaction of electromagnetic waves
and matter. The scattering of sun’s rays by raindrops to produce a rainbow and
appearance of a colorful spectrum when a narrow beam of sunlight is passed
through a triangular glass prism are the simple examples where white light is
separated into the visible spectrum of primary colors. This visible light is
merely a part of the whole spectrum of electromagnetic radiation, extending
from the radio waves to cosmic rays. All these apparently different forms of
electromagnetic radiations travel at the same velocity but characteristically
differ from each other in terms of frequencies and wavelength.
Absorption of
Different Electromagnetic radiations
In absorption spectroscopy, though the
mechanism of absorption of energy is different in the ultraviolet, infrared and
nuclear magnetic resonance regions, the fundamental process is the absorption
of a discrete amount of energy. The energy required for the transition from a
state of lower energy (E1) to state of higher energy (E2) is exactly equivalent
to the energy of electromagnetic radiation that causes transition.
Fig.
Energy transition for the absorption of any electromagnetic radiation
E
. .. 1 – E2 = E = hν = h c / λ
Where
E is energy of electromagnetic radiation being absorbed, h is the universal
Planck’s constant, 6.624 x 10-27 erg sec and ν is the frequency of incident
light in cycles per second (cps or hertz, Hz), c is velocity of light 2.998 x
1010 cm s -1 and λ = wavelength (cm)
Therefore,
higher is the frequency, higher would be the energy and longer is the
wavelength, lower would be the energy. As we move from cosmic radiations to
ultraviolet region to infrared region and then radio frequencies, we are
gradually moving to regions of lower energies.
A
molecule can only absorb a particular frequency, if there exists within the
molecule an energy transition of magnitude E = h ν
Although
almost all parts of electromagnetic spectrum are used for understanding the
matter, in organic chemistry we are mainly concerned with energy absorption
from only ultraviolet and visible, infrared, microwave and radiofrequency
regions.
Ultraviolet –
visible spectroscopy (λ 200 - 800 nm) studies the changes
in electronic energy levels within the molecule arising due to transfer of
electrons from π- or non-bonding orbitals. It commonly provides the knowledge
about π-electron systems, conjugated unsaturations, aromatic compounds and
conjugated non-bonding electron systems etc.
Infrared
spectroscopy ( ν 400-4000 cm-1) studies the
changes in the vibrational and rotation movements of the molecules. It is
commonly used to show the presence or absence of functional groups which have
specific vibration frequencies viz. C=O, NH2, OH, CH, C-O etc.
In
the present chapter, UV-Vis and Infrared spectroscopy have been discussed.
Ultraviolet
and Visible Spectroscopy This
absorption spectroscopy uses electromagnetic radiations between 190 nm to 800
nm and is divided into the ultraviolet (UV, 190-400 nm) and visible (VIS,
400-800 nm) regions. Since the absorption of ultraviolet or visible radiation
by a molecule leads transition among electronic energy levels of the molecule,
it is also often called as electronic spectroscopy. The information provided by
this spectroscopy when combined with the information provided by NMR and IR
spectral data leads to valuable structural proposals.
Principles
of Absorption Spectroscopy : Beer’s and Lambert’s Law The greater the number of
molecules that absorb light of a given wavelength, the greater the extent of
light absorption and higher the peak intensity in absorption spectrum. If there
are only a few molecules that absorb radiation, the total absorption of energy
is less and consequently lower intensity peak is observed. This makes the basis
of Beer-Lambert Law which states that the fraction of incident radiation
absorbed is proportional to the number of absorbing molecules in its path.
When
the radiation passes through a solution, the amount of light absorbed or
transmitted is an exponential function of the molecular concentration of the
solute and also a function of length of the path of radiation through the
sample.
Therefore,
Log Io / I = ε c l
Where
Io = Intensity of the incident light (or the light intensity passing through a
reference cell)
I = Intensity of light transmitted through the
sample solution
c = concentration of the solute in mol l-1
l
= path length of the sample in cm
ε
= molar absorptivity or the molar extinction coefficient of the substance whose
light absorption is under investigation. It is a constant and is a
characteristic of a given absorbing species (molecule or ion) in a particular
solvent at a particular wavelength. ε is numerically equal to the absorbance of
a solution of unit molar concentration (c = 1) in a cell of unit length ( l =
1) and its units are liters.moles-1 . cm -1. However, it is customary practice
among organic chemists to omit the units.
The
ratio I / Io is known as transmittance T and the logarithm of the inverse ratio
Io / I is known as the absorbance A.
-
Log I / Io = - log T = ε c l
and
Log Io / I = A = ε c l
or
A = ε c l
For
presenting the absorption characteristics of a spectrum, the positions of peaks
are reported as λmax (in nm) values and the absorptivity is expressed in
parenthesis.
Solvent
Effects
Highly pure, non-polar solvents such as
saturated hydrocarbons do not interact with solute molecules either in the
ground or excited state and the absorption spectrum of a compound in these
solvents is similar to the one in a pure gaseous state. However, polar solvents
such as water, alcohols etc. may stabilize or destabilize the molecular
orbitals of a molecule either in the ground state or in excited state and the
spectrum of a compound in these solvents may significantly vary from the one
recorded in a hydrocarbon solvent.
(i)
π -π*
Transitions
In case of π Æ π* transitions,
the excited states are more polar than the ground state and the dipole-dipole
interactions with solvent molecules lower the energy of the excited state more
than that of the ground state. Therefore a polar solvent decreases the energy
of π Æ π* transition and absorption maximum appears ~10-20 nm red shifted in
going from hexane to ethanol solvent.
(ii)
n -π*
Transitions
In case of n Æ π* transitions, the polar
solvents form hydrogen bonds with the ground state of polar molecules more
readily than with their excited states. Therefore, in polar solvents the
energies of electronic transitions are increased. For example, the figure 5
shows that the absorption maximum of acetone in hexane appears at 279 nm which
in water is shifted to 264 nm, with a blue shift of 15 nm.
Fig: UV-spectra of acetone in
hexane and in water
