Lecture 1
Spectroscopy
• It is the branch of science that deals with the
study of interaction of electromagnetic
radiation with matter.
Electromagnetic Radiation
• Electromagnetic radiation consist of discrete
packages of energy which are called as
photons.
• A photon consists of an oscillating electric field
(E) & an oscillating magnetic field (M) which
are perpendicular to each other.
Electromagnetic Radiation
• Frequency (ν):
– It is defined as the number of times electrical field
radiation oscillates in one second.
– The unit for frequency is Hertz (Hz).
1 Hz = 1 cycle per second
• Wavelength (λ):
– It is the distance between two nearest parts of the
wave in the same phase i.e. distance between two
nearest crest or troughs.
Electromagnetic Radiation
• The relationship between wavelength &
frequency can be written as:
c=νλ
• As photon is subjected to energy, so
E=hν=hc/λ
Visible light:
• The visible spectrum is the electromagnetic
spectrum that is visible to the human eye.
• The longest wavelength is red and the shortest
is violet.
Electromagnetic Radiation
When visible light is passed through a prism, it split up into 7 coloures which
correspond to definite wavelengths. This phenomenon is called dispersion.
Violet 400 - 420 nm
Yellow 570 - 585 nm
Indigo 420 - 440 nm
Orange 585 - 620 nm
Blue 440 - 490 nm
Red 620 - 780 nm
Green 490 - 570 nm
Electromagnetic spectrum:
• The visible spectrum constitutes but a small
part of the total radiation spectrum. Most of
the radiation that surrounds us cannot be
seen, but can be detected by dedicated
sensing instruments. This electromagnetic
spectrum ranges from very short wavelengths
(including gamma and x-rays) to very long
wavelengths (including microwaves and
broadcast radio waves). The following chart
displays many of the important regions of this
spectrum.
Electromagnetic Radiation
Principles of Spectroscopy
• The principle is based on the measurement of
spectrum of a sample containing atoms /
molecules.
• Spectrum is a graph of intensity of absorbed or
emitted radiation by sample verses frequency
(ν) or wavelength (λ).
• Spectrometer is an instrument design to
measure the spectrum of a compound.
Principles of Spectroscopy
1. Absorption Spectroscopy:
• An analytical technique which concerns with
the
measurement
of
absorption
of
electromagnetic radiation.
• e.g. UV (185 - 400 nm) / Visible (400 - 800 nm)
Spectroscopy, IR Spectroscopy (0.76 - 15 μm)
Principles of Spectroscopy
2. Emission Spectroscopy:
• An analytical technique in which emission
(of a particle or radiation) is dispersed
according to some property of the emission
& the amount of dispersion is measured.
• e.g. Mass Spectroscopy
UV-Visible Spectroscopy
• A diagram of the components of a typical spectrometer are shown
in the following diagram. The functioning of this instrument is
relatively straightforward. A beam of light from a visible and/or UV
light source (colored red) is separated into its component
wavelengths by a prism or diffraction grating. Each monochromatic
(single wavelength) beam in turn is split into two equal intensity
beams by a half-mirrored device. One beam, the sample beam
(colored magenta), passes through a small transparent container
(cuvette) containing a solution of the compound being studied in a
transparent solvent. The other beam, the reference (colored blue),
passes through an identical cuvette containing only the solvent. The
intensities of these light beams are then measured by electronic
detectors and compared. The intensity of the reference beam,
which should have suffered little or no light absorption, is defined
as I0. The intensity of the sample beam is defined as I. Over a short
period of time, the spectrometer automatically scans all the
component wavelengths in the manner described. The ultraviolet
(UV) region scanned is normally from 200 to 400 nm, and the visible
portion is from 400 to 800 nm.
• If the sample compound does not absorb light of of a given wavelength, I =
I0. However, if the sample compound absorbs light then I is less than I0,
and this difference may be plotted on a graph versus wavelength, as
shown on the right. Absorption may be presented as transmittance (T =
I/I0) or absorbance (A= log I0/I). If no absorption has occurred, T = 1.0 and
A= 0. Most spectrometers display absorbance on the vertical axis, and the
commonly observed range is from 0 (100% transmittance) to 2 (1%
transmittance). The wavelength of maximum absorbance is a
characteristic value, designated as λmax. Different compounds may have
very different absorption maxima and absorbances. Intensely
absorbing compounds must be examined in dilute solution, so that
significant light energy is received by the detector, and this requires
the use of completely transparent (non-absorbing) solvents. The
most commonly used solvents are water, ethanol, hexane and
cyclohexane. Solvents having double or triple bonds, or heavy
atoms (e.g. S, Br & I) are generally avoided. Because the absorbance
of a sample will be proportional to its molar concentration in the
sample cuvette, a corrected absorption value known as the molar
absorptivity is used when comparing the spectra of different
compounds. This is defined as:
• Molar Absorptivity,ε = A/ c l
( where A= absorbance, c = sample concentration in moles/liter & l =
length of light path through the cuvette in cm.)
• For the spectrum on the right, a solution of 0.249 mg of the
unsaturated aldehyde in 95% ethanol (1.42 • 10-5 M) was placed in
a 1 cm cuvette for measurement. Using the above formula, ε =
36,600 for the 395 nm peak, and 14,000 for the 255 nm peak. Note
that the absorption extends into the visible region of the spectrum,
so it is not surprising that this compound is orange colored.
Molar absorptivities may be very large for strongly absorbing
compounds (ε >10,000) and very small if absorption is weak (ε =
10 to 100).
UV-Visible Absorption Spectra
• To understand why some compounds are colored and others are
not, and to determine the relationship of conjugation to color, we
must make accurate measurements of light absorption at different
wavelengths in and near the visible part of the spectrum.
Commercial optical spectrometers enable such experiments to be
conducted with ease, and usually survey both the near ultraviolet
and visible portions of the spectrum. The visible region of the
spectrum comprises photon energies of 36 to 72 kcal/mole, and the
near ultraviolet region, out to 200 nm, extends this energy range to
143 kcal/mole. Ultraviolet radiation having wavelengths less than
200 nm is difficult to handle, and is seldom used as a routine tool
for structural analysis. The energies noted above are sufficient to
promote or excite a molecular electron to a higher energy orbital.
• Consequently, absorption spectroscopy carried out in this region is
sometimes called "electronic spectroscopy". A diagram showing the
various kinds of electronic excitation that may occur in organic
molecules is shown on the left. Of the six transitions outlined, only
the two lowest energy ones (left-most, colored blue) are achieved
by the energies available in the 200 to 800 nm spectrum. As a rule,
energetically favored electron promotion will be from the highest
occupied molecular orbital (HOMO) to the lowest unoccupied
molecular orbital (LUMO), and the resulting species is called an
excited state.
TYPES OF ELECTRONIC TRANSITIONS
FOUR TYPES OF ELECTRONIC
TRANSITIONS
1
σ→σ* for an ordinary carbon-carbon bond. (125-150)nm
2
π→π* for an isolated double bond. (160-190)nm
3 compounds. Aldehydes and ketones(275-295)nm
n→π* In carbonyl
4
n→σ* in oxygen,nitrogen
,sulphur and Halogen compounds.
(150-250)nm
21
When sample molecules are exposed to light having an
energy that matches a possible electronic transition within
the molecule, some of the light energy will be absorbed as
the electron is promoted to a higher energy orbital. An
optical spectrometer records the wavelengths at which
absorption occurs, together with the degree of absorption
at each wavelength. The resulting spectrum is presented as
a graph of absorbance (A) versus wavelength, as in the
isoprene spectrum shown below. Since isoprene is
colorless, it does not absorb in the visible part of the
spectrum and this region is not displayed on the graph.
Absorbance usually ranges from 0 (no absorption) to 2
(99% absorption), and is precisely defined in context with
spectrometer operation.
Because the absorbance of a sample will be proportional to the
number of absorbing molecules in the spectrometer light beam
(e.g. their molar concentration in the sample tube), it is necessary
to correct the absorbance value for this and other operational
factors if the spectra of different compounds are to be compared
in a meaningful way. The corrected absorption value is called
"molar absorptivity", and is particularly useful when comparing
the spectra of different compounds and determining the relative
strength of light absorbing functions (chromophores). The functional
groups containing multiple bonds capable of absorbing radiations above 200
nm due to n → π* & π → π* transitions.
e.g. NO2, N=O, C=O, C=N, C≡N, C=C, C=S, etc
• If the isoprene spectrum on the right was obtained from a dilute
hexane solution (c = 4 * 10-5 moles per liter) in a 1 cm sample
cuvette, a simple calculation using the above formula indicates a
molar absorptivity of 20,000 at the maximum absorption
wavelength. Indeed the entire vertical absorbance scale may be
changed to a molar absorptivity scale once this information about
the sample is in hand. Clicking on the spectrum will display this
change in units. From the chart below it should be clear that the
only molecular moieties likely to absorb light in the 200 to 800 nm
region are pi-electron functions and hetero atoms having nonbonding valence-shell electron pairs. Such light absorbing groups
are referred to as chromophores. A list of some simple
chromophores and their light absorption characteristics is provided
on the left above. The oxygen non-bonding electrons in alcohols
and ethers do not give rise to absorption above 160 nm.
Consequently, pure alcohol and ether solvents may be used for
spectroscopic studies.
• The presence of chromophores in a molecule is best documented
by UV-Visible spectroscopy, but the failure of most instruments to
provide absorption data for wavelengths below 200 nm makes the
detection of isolated chromophores problematic. Fortunately,
conjugation generally moves the absorption maxima to longer
wavelengths, as in the case of isoprene, so conjugation becomes
the major structural feature identified by this technique.
Molar absorptivities may be very large for strongly absorbing
chromophores (>10,000) and very small if absorption is weak (10 to
100). The magnitude ofε reflects both the size of the chromophore
and the probability that light of a given wavelength will be absorbed
when it strikes the chromophore.
Chromophore
Example
Excitation
λmax, nm
ε
Solvent
C=C
Ethene
π
__>
π*
171
15,000
hexane
C≡C
1-Hexyne
π
__>
π*
180
10,000
hexane
C=O
Ethanal
n
π
__>
π*
__> π*
290
180
15
10,000
hexane
hexane
N=O
Nitromethane
n
π
__>
π*
π*
275
200
17
5,000
ethanol
ethanol
C-X X=Br
X=I
Methyl bromide
Methyl Iodide
n
n
__>
σ*
σ*
205
255
200
360
hexane
hexane
__>
__>
Applications
• Qualitative & Quantitative Analysis:
– It is used for characterizing aromatic compounds
and conjugated olefins.
– It can be used to find out molar concentration of the
solute under study.
• Detection of impurities:
– It is one of the important method to detect
impurities in organic solvents.
• Detection of isomers are possible.
• Determination of molecular weight using Beer’s
law.