MODULE IV

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 (Table 1).

The propagation of these radiations involves both electric and magnetic forces which give rise to
their common class name electromagnetic radiation. In spectroscopy, only the effects associated with

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electric component of electromagnetic wave are important. Therefore, the light wave traveling through
space is represented by a sinusoidal trace (figure 1). In this diagram λ is the wavelength and distance A is
known as the maximum amplitude of the wave. Although a wave is frequently characterized in terms of
its wavelength λ, often the terms such as wavenumber , frequency (ν), cycles per second (cps) or hertz
(Hz) are also used.

Figure 1: Wave like propagation of light ( λ = wavelength , A = amplitude)

The unit commonly used to describe the wavelength is centimeters (cm), the different units are
used to express the wavelengths in different parts of the electromagnetic spectrum. For example, in the
ultraviolet and visible region, the units use are angstrom (Ǻ) and nanometer (nm). In the infrared region,
the commonly used unit is wavenumber (cm-1), which gives the number of waves per centimeter. The four
quantities wavelength, wavenumber, frequency and velocity can be related to each other by following
relationships:

Emission and Absorption spectroscopy

When energy in the form of light, heat, or chemical agents is given to an element, the electrons
of its atoms accept the energy and go to higher energy levels. However, these electrons have to emit
energy in order to return to their ground state, since the excited state is unstable. The frequencies of light
emitted in such a case constitute the emission spectrum (Figure 2). When an electron comes down from
an excited state to the ground state, it emits a photon of energy. The energy of this photon depends on
the difference between the energy levels of the excited state and ground state of that electron.

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Figure 2: Concept of Emission spectrum

The intensity of the light emitted (at the specific frequency corresponding to the energy level
difference) is proportional to concentration of the analyte. Spectroscopic studies based of this emitted
light is known as emission spectroscopy.

Flame emission spectroscopy or Flame photometry is a common example of application of

emission spectroscopy. However, Flame emission spectroscopy is limited only to alkali and alkaline earth
elements as the ionization energy of only these elements are low enough to be thermally excited to yield
a emission spectrum (Figure 3).

Figure 3: Flame emission spectroscopy or Flame photometry

Inductively Coupled Plasma –Atomic emission spectroscopy (ICP-AES) spectroscopy is a superior

example that involves injection of the sample into the plasma, resulting in emission spectrum for all
elements present and subsequently, their estimation (Figure 4).

Although each element emits energy at multiple wavelengths, in the ICP-AES technique it is most
common to select a single wavelength (or a very few) for a given element. The intensity of the energy
emitted at the chosen wavelength is proportional to the amount (concentration) of that element in the
sample being analyzed. Thus, by determining which wavelengths are emitted by a sample and by
determining their intensities, the analyst can qualitatively and quantitatively find the elements from the
given sample relative to a reference standard.

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 Figure 4: Inductively Coupled Plasma Atomic Emission Spectroscopy (ICP-AES)

Electrons of an element which are in the ground state may absorb incident energy in order to
reach a higher energy state. The frequencies of light transmitted through this substance, with dark bands
showing absorbed light, constitute the absorption spectrum of the substance.

Figure 5: The concept of Absorption spectroscopy

The spectroscopic studies based on ‘absorbance’ of light at concerned frequency (corresponding

to the energy gap) is called absorption spectroscopy (Figure 5).

All types of spectroscopy discussed heron are absorption spectroscopy.

Lambert – Beer’s law

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Lambert's law stated that absorbance of a material sample is directly proportional to its thickness (path
length). Much later, August Beer discovered another attenuation relation in 1852. Beer's law stated that
absorbance is proportional to the concentrations of the attenuating species in the material sample.

Combination of these two laws resulted in Lambert – beer’s law, which states that absorbance of a
monochromatic light passing through a medium is proportional to both thickness(l) of the medium as well
as the concentration (c) of the absorbing species.

Mathematically,

Figure 6: Concept of Lambert-Beer’s Law

Limitations of Lambert – Beer’s law

The linearity of the Beer-Lambert law is limited by chemical and instrumental factors. Causes of
nonlinearity include:

• deviations in absorptivity coefficients at high concentrations (>0.01M) due to electrostatic

interactions between molecules in close proximity
• scattering of light due to particulates in the sample
• fluoresecence or phosphorescence of the sample
• changes in refractive index at high analyte concentration
• shifts in chemical equilibria as a function of concentration
• non-monochromatic radiation, deviations can be minimized by using a relatively flat part of the
absorption spectrum such as the maximum of an absorption band
• stray light

UV – VIS Spectroscopy

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 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. Nature of Electronic Transitions The total energy of a molecule is the sum of its electronic, its
vibrational energy and its rotational energy. Energy absorbed in the UV region produces changes in the
electronic energy of the molecule. As a molecule absorbs energy, an electron is promoted from an
occupied molecular orbital (usually a non-bonding n or bonding π orbital) to an unoccupied molecular
orbital (an antibonding π∗ or σ* orbital) of greater potential energy (figure 3). For most molecules, the
lowest-energy occupied molecular orbitals are σ orbitals, which correspond to σ bonds. The π orbitals lie
at relatively higher energy levels than σ orbitals and the non-bonding orbitals that hold unshared pairs of
electrons lie even at higher energies. The antibonding orbitals (π* and σ * ) are orbitals of highest energy.
The relative potential energies of these orbitals and vaious possible transitions have been depicted in
figure 7.

Figure 7: Relative energies of orbitals most commonly involved in electronic spectroscopy of organic
molecules

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The saturated aliphatic hydrocarbons (alkanes) exhibit only σ → σ* transitions but depending on the
functional groups the organic molecules may undergo several possible transitions which *can be placed
in the increasing order of their energies viz. n → π* < n → σ * < π→ π* < σ → π* < σ → σ* . Since all these
transitions require fixed amount of energy (quantized), an ultraviolet or visible spectrum of a compound
would consist of one or more well defined peaks, each corresponding to the transfer of an electron from
one electronic level to another. If the differences between electronic energy levels of two electronic states
are well defined i.e. if the nuclei of the two atoms of a diatomic molecule are held in fixed position, the
peaks accordingly should be sharp. However, vibrations and rotations of nuclei occur constantly and as a
result each electronic state in a molecule is associated with a large number of vibrational and rotational
states. At room temperature, the molecules in the ground state will be in the zero vibrational level (Gυ o).
This is shown schematically in figure 8.

Figure 8: Energy level diagram showing excitation between different vibrational and rotational levels of
two electronic states

The transition of an electron from one energy level to another is thus accompanied by
simultaneous change in vibrational and rotational states and causes transitions between various
vibrational and rotational levels of lower and higher energy electronic states. Therefore many radiations
of closely placed frequencies are absorbed and a broad absorption band is obtained. When a molecule
absorbs ultraviolet or visible light of a defined energy, an assumption is made that only one electron is

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excited form bonding orbital or non-bonding orbital to an anti-bonding orbital and all other electrons
remain unaffected. The excited state thus produced is formed in a very short time i.e. of the order of 10-
15 seconds. In accordance with Franck-Condon principle, during electronic excitation the atoms of the
molecule do not move. The most probable transition would appear to involve the promotion of one
electron from the highest occupied molecular orbital (HOMO) to the lowest unoccupied molecular orbital
(LUMO), but in many cases several transitions can be observed, giving several absorption bands in the
spectrum. We can have a general view of the possible transitions prevailing in organic compounds.

Alkanes can only undergo σ → σ* transitions. These are high-energy transitions and involve very short
wavelength ultraviolet light (< 150 nm). These transitions usually fall out-side the generally available
measurable range of UV-visible spectrophotometers (200-1000 nm). The σ → σ* transitions of methane
and ethane are at 122 and 135 nm, respectively. In alkenes amongst the available σ → σ* and π → π*
transitions, the π → π* transitions are of lowest energy and absorb radiations between 170-190 nm. In
saturated aliphatic ketones the lowest energy transition involves the transfer of one electron of the
nonbonding electrons of oxygen to the relatively low-lying π* anti-bonding orbital. This n → π* transition
is of lowest energy (~280 nm) but is of low intensity as it is symmetry forbidden. Two other available
transitions are n → π* and π → π*. The most intense band for these compounds is always due to π → π*
transition. In conjugated dienes the π → π* orbitals of the two alkene groups combine to form new
orbitals – two bonding orbitals named as π1 and π2 and two antibonding orbitals named as π3* and π4*.
It is apparent that a new π → π* transition of low energy is available as a result of conjugation. Conjugated
dienes as a result absorb at relatively longer wavelength than do isolated alkenes (see figure 10 ).

Spectral Measurements

The UV-Vis spectra are usually measured in very dilute solutions and the most important criterion in the
choice of solvent is that the solvent must be transparent within the wavelength range being examined.
Table 2 lists some common solvents with their lower wavelength cut off limits. Below these limits, the
solvents show excessive absorbance and should not be used to determine UV spectrum of a sample.

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 → 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

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of π → π* transition and absorption maximum appears ~10-20 nm red shifted in going from hexane to
ethanol solvent.

Table 2:

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(ii) n→ π* Transitions In case of n → 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 9 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 15nm.

Figure 9 : UV-spectra of acetone in hexane and in water

Some important terms and definitions

(i) Chromophore: The energy of radiation being absorbed during excitation of electrons from ground
state to excited state primarily depends on the nuclei that hold the electrons together in a bond. The
group of atoms containing electrons responsible for the absorption is called chromophore. Most of the
simple un-conjugated chromophores give rise to high energy transitions of little use. Some of these
transitions have been listed in table 3.

Table 3:

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 For example, alkanes contain only single bonds with only possible σ→σ* type electronic
transitions. These transitions absorb radiations shorter than wavelengths that are experimentally
accessible in usually available spectrophotometers. In saturated molecules with heteroatom bearing non-
bonding pairs of electrons, n→σ* transitions become available. These are also high energy transitions. In
unsaturated compounds, π→π* transitions become possible. Alkenes and alkynes absorb ~ 170 nm but
the presence of substituents significantly affects their position. The carbonyl compounds and imines can
also undergo n →π* transitions in addition to π →π*. Amongst these, the most studied transitions are
n→π* as these absorb at relatively longer wavelength 280-300 nm. These are low intensity (ε 10 -100)
transitions.

(ii) Auxochrome: The substituents that themselves do not absorb ultraviolet radiations but their presence
shifts the absorption maximum to longer wavelength are called auxochromes. The substituents like
methyl, hydroxyl, alkoxy, halogen, amino group etc. are some examples of auxochromes.

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(iii) Bathochromic Shift or Red shift: A shift of an absorption maximum towards longer wavelength or
lower energy.

(iv) Hypsochromic Shift or Blue Shift: A shift of an absorption maximum towards shorter wavelength or
higher energy.

(v) Hypochromic Effect: An effect that results in decreased absorption intensity. (vi) Hyperchromic Effect:
An effect that results in increased absorption intensity.

Applications of Electronic Spectroscopy in Predicting Absorption Maxima of Organic Molecules

1: Conjugated Dienes, Trienes and Polyenes

The presence of conjugate double bond decreases the energy difference between HOMO and
LUMO of resulting diene. The figure 10 shows the change in energy of MO on conjugation. As a result, the
radiations of longer wavelength are absorbed. The conjugation not only results in bathochromic shift
(longer wavelength) but also increases the intensity of absorption. As the number of conjugated double
bonds is increased, the gap between highest occupied molecular orbital (HOMO) and lowest unoccupied
molecular orbital (LUMO) is progressively lowered. Therefore, the increase in size of the conjugated
system gradually shifts the absorption maximum (λmax) to longer wavelength and also increases the
absorption. For example, ethylene absorbs at 175 nm (ε = 1000) and the conjugation in butadiene gives a
strong absorption at longer wavelength at 230 nm and with higher intensity (ε = >1000).

Figure 10: Effect of conjugation on orbitals’ energy and respective π-π* transition

The presence of alkyl substituents on double bond also produces bathochromic shift and
hyperchromic effect. These effects are additive in dienes and up to some extent in trienes. The open chain
dienes can achieve s-cis or s-trans conformations and similarly diene system can be homoannular or
heteroannular in cyclic systems. In 1941, Woodward suggested empirical rules for predicting the
absorption of open chain and six-membered ring dienes which have been later on extended to large
number of dienes and trienes (Table 4 ).

Table 4:

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 For example, here the absorption maxima for dienes 1 and 2 have been calculated according to
Woodward rules. The comparison of calculated λmax values with observed λmax values highlights the
importance of these rules.

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 As the number of double bonds in conjugation increases, the bathochromic (towards longer
wavelength) shift in lowest energy absorption maxima is observed. The increase in conjugation gradually
shifts the maxima to visible region (> 400 nm) and imparts colour to the sample. Table 5 shows the λmax
shift in Me(CH=CH)nMe with increasing number of conjugated double bonds. β - Carotene (figure 11)
responsible for red color in carrots is a typical example of polyene with 11 conjugated double bonds and
exhibits λmax at 445 nm.

Figure 11: Structure of β-carotene

Table 5:

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 2: Carbonyl Compounds

Carbonyl compounds have two principal UV radiations, the allowed π →π* transitions and the
forbidden n →π* transitions. In amides, acids, esters or acid halides, the substituents viz. NR 2, OH, OR, or
–X on carbonyl group show pronounced hypsochromic effect on the n →π* transitions. The hypsochromic
effect is due to inductive effect of nitrogen, oxygen or halogen atoms. The heteroatom withdraws
electrons from carbonyl carbon and makes carbonyl oxygen lone pair of electrons more stabilized due to
its involvement in increasing C=O bond order. As a result, the n → π* transition of these compounds is
shifted to 200-215 nm range relative to 270 nm in aldehydes and ketones. Conjugation of the carbonyl
group with double bond shifts both n →π* and π →π* transitions to longer wavelengths. The effect on
π→ π* band is more pronounced. Woodward formulated rules to predict the position of an absorption
maximum in an unknown enone. These rules have been summarized in table 6.

Table 6:

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 3. Aromatic Compounds

The simplest aromatic compound is benzene. It shows two primary bands at 184 (ε = 47,000) and
202 (ε = 7400) nm and a secondary fine structure band at 255 nm (ε = 230 in cyclohexane). Substituents
on the benzene ring also cause bathochromic and hypsochromic shifts of various peaks. Unlike dienes and
unsaturated ketones, the effects of various substituents on the benzene ring are not predictable.
However, qualitative understanding of the effects of substituents on the characteristics of UV-Vis
spectrum can be considered by classifying the substituents into electron-donating and electron-
withdrawing groups. (i) Effect of Substituents with Unshared Electrons: The non-bonding electrons
increase the length of π-system through resonance and shift the primary and secondary absorption bands
to longer wavelength. More is the availability of these non-bonding electrons, greater the shift will be. In

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addition, the presence of non-bonding electrons introduces the possibility of n → π* transitions. If non-
bonding electron is excited into the extended π*chromophore, the atom from which it is removed
becomes electron-deficient and the π-system of aromatic ring becomes electron rich. This situation causes
a separation of charge in the molecule and such excited state is called a charge-transfer or an electron-
transfer excited state. In going from benzene to t-butylphenol, the primary absorption band at 203.5 nm
shifts to 220 nm and secondary absorption band at 254 nm shifts to 275 nm. Further, the increased
availability of n electrons in negatively charged t-butylphenoxide ion shifts the primary band from 203.5
to 236 nm (a 32.5 nm shift) and secondary band shifts from 254 nm to 290 nm (a 36 nm shift) (Figure 12
). Both bands show hyperchromic effect. On the other hand, in the case of anilinium cation, there are no
n electrons for interaction and absorption properties are quite close to benzene. But in aniline, the
primary band is shifted to 232 nm from 204 nm in anilinium cation and the secondary band is shifted to
285 nm from 254 nm (Figure 13).

Figure 12 : UV-spectra of t-butyl phenol and t-buty phenoxide in methanol

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 Figure 13 : UV-spectra of aniline and anilinium salt in methanol

(ii) Effect of π Conjugation: Conjugation of the benzene ring also shifts the primary band at 203.5
nm more effectively to longer wavelength and secondary band at 254 nm is shifted to longer wavelength
to lesser extent. In some cases, the primary band overtakes the secondary band. For example, benzoic
acid shows primary band at 250 nm and secondary band at 273 nm, but cinnamic acid that has longer
chromophore exhibits primary band at 273 nm and secondary band remains merged with it. Similarly, in
benzaldehyde, the secondary band appears at 282 nm and primary band at 242 nm but in case of
cinnamaldehyde, primary band appears at 281 nm and remains merged with secondary band (figure 14).
The hyperchromic effect arising due to extended conjugation is also visible.

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 Figure 14 : UV-spectra of benzaldehyde and cinnamaldehyde in methanol

(iii) Effect of Electron-withdrawing and Electron-releasing Groups: Electron-withdrawing

substituents viz. NH3 + , SO2NH2, CN, COOH, COCH3, CHO and NO2 etc. have no effect on the position of
secondary absorption band of benzene ring. But their conjugation effects with πelectrons of the aromatic
ring are observed. Electron-donating groups such as -CH3, -Cl, -Br, - OH, -OCH3, -NH2 etc increase both λmax
and εmax values of the secondary band.

In case of disubstituted benzene derivatives, it is essential to consider the effect of both the
substituents. In para-substituted benzenes, two possibilities exist. If both the groups are electron-
donating then the observed spectrum is closer to monosubstituted benzene. The group with stronger
effect determines the extent of shifting of primary band. If one group is electron-releasing and other is
electron-withdrawing, the magnitude of red shift is grater compared to the effect of single substituent
individually. This is attributed to the increased electron drift from electron-donating group to the electron-
withdrawing group through π-bond of benzene ring. For example, aniline shows secondary band at 285
nm which due to presence of electron-withdrawing p-nitro substituent is shifted to 367 nm with a
significant increase in absorptivit (figure 15).

Figure 15: UV-spectra of aniline and p-nitroaniline in methanol

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 If two groups of a disubstituted benzene derivative are placed ortho- or meta- to each other, the
combined effect of two substiuents is observed. In case of substituted benzoyl derivatives, an empirical
correction of structure with observed position of the primary absorption band has been developed. In the
absence of steric hindrance to co-planarity, the calculated values are within ± 5 nm of the observed value.

(iv) Polycyclic Aromatic Compounds: In case of polycyclic aromatic hydrocarbons, due to

extended conjugation, both primary and secondary bands are shifted to longer wavelength. These spectra
are usually complicated but are characteristic of parent compound. The primary band at 184 nm in
benzene shifts to 220 nm in case of naphthalene and 260 nm in case of anthracene.

Similarly, the structured secondary band which appears as broad band around 255 nm in benzene
is shifted to 270 nm and 340 nm respectively in case of naphthalene and anthracene molecules.

Commercial Applications of UV and Visible Spectroscopy

The UV-Vis spectroscopy has innumerable applications in the drugs and pharmaceutical industry.
Beer-Lambert law offers a valuable and simple method for quantitative analysis. In practice, a calibration
curve is constructed by plotting absorbance vs. molar concentration and the concentration of unknown
with ‘X’ absorbance is determined by finding the concentration corresponding to the measured
absorbance on the calibration curve. The UV spectroscopy is used extensively in determining rate
constants, equilibrium constants, acid-base dissociation constants etc for chemical reactions. The use of
UV spectrometry in evaluation of enzymatic assays has become very common e.g. the activity of enzyme
dehydrase is assayed by measuring the formation of ergosterol at 282 nm.

Infrared Absorption Spectroscopy (Vibrational spectroscopy)

The two atoms joined together by a chemical bond (may be single, double or triple bond),
macroscopically can be composed as two balls joined by a spring. The application of a force like (i)
stretching of one or both the balls (atoms) away from each other or closer to each other (ii) bending of
one of the atoms either vertically or horizontally and then release of the force results in the vibrations on
the two balls (atoms). These vibrations depend on the strength of the spring and also the mode (stretching
or bending) in which the force is being applied. Similarly, at ordinary temperatures, organic molecules are
in a constant state of vibrations, each bond having its characteristic stretching and bending frequencies.
When infrared light radiations between 4000-400 cm-1 (the region most concerned to an organic chemist)
are passed through a sample of an organic compound, some of these radiations are absorbed by the
sample and are converted into energy of molecular vibrations. The other radiations which do not interact
with the sample are transmitted through the sample without being absorbed. The plot of % transmittance
against frequency is called the infrared spectrum of the sample or compound. This study of vibrations of
bonds between different atoms and varied multiplicities which depending on the elctronegativity, masses
of the atom and their geometry vibrate at different but specified frequencies; is called infrared

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spectroscopy. The presence of such characteristic vibrational bands in an infrared spectrum indicates the
presence of these bonds in the sample under investigation.

Hooke’s law and Absorption of radiations The band positions in the IR spectrum are presented in
wave numbers ( ν ) whose unit is the reciprocal centimeter (cm-1). ν is proportional to the energy of
vibration. ∆E = hυ = hc / λ = hc ν Therefore, in principle, each absorption of radiation in the infrared region
is quantized and should appear as sharp line. However, each vibrational transition within the molecule is
associated with number of rotational energy changes and thus appears as combination of vibrational-
rotational bands. The analogy of a chemical bond with two atoms linked through a spring can be used to
rationalize several features of the infrared spectroscopy. The approximation to vibration frequency of a
bond can be made by the application of Hooke’s law. In Hooke’s law, two atoms and their connecting
bond are treated as a simple harmonic oscillator composed of two masses joined by a spring and
frequency of vibration is stated as

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 Therefore, the vibrational frequency of a bond would increase with the decrease in reduced mass
of the system. It implies that C-H and O-H stretching absorptions should appear at higher frequencies than
C-C and C-O stretching frequencies. Similarly, O-H stretching should appear at higher frequency than O-D
stretching. Further, in parallel with the general knowledge that the stretching of the spring requires more
energy than to bend it, the stretching absorption of a band always appear at higher energy than the
bending absorption of the same band. The Hooke’s law can be used to theoretically calculate the
approximate stretching frequency of a bond. The value of K is approximately 5x10 5 dyne/cm for single
bonds and approximately two and three times this value for the double and triple bonds, respectively.

Let us calculate the approximate frequency of the C-H starching vibration from the masses of
carbon and hydrogen

mC = mass of carbon atom = 20x10-24 g

mH = mass of hydrogen atom = 1.6x10-24 g

Let us consider how the radiations are being absorbed. We know that at ordinary temperature,
molecules are in constant state of vibrations. The change in dipole moment during vibration of the
molecule produces a stationary alternating electric field. When the frequency of incident electromagnetic
radiations is equal to the alternating electric field produced by changes in the dipole moment, the
radiation is absorbed and vibrational levels of the molecule are excited. Once in the vibrationally excited
state, the molecules can loose the extra energy by rotational, collision or translational processes etc. and
come back to ground state. Therefore, only those vibrations which result in a rhythmical change in the
dipole moment of the molecule absorb infrared radiations and are defined as IR active. The others which
do not undergo change in dipole moment of the molecule are IR inactive e.g. the stretching of a
symmetrically substituted bond, viz ‘carbon triple bond carbon’ in acetylene and symmetrical stretching
in carbon dioxide (figure 17) – a linear molecule, produce no change in the dipole moment of the system
and these vibrations cannot interact with infrared light and are IR inactive. In general, the functional
groups that have a strong dipole give rise to strong absorption bands in the IR.

Modes of molecular vibrations

Molecules with large number of atoms possess a large number of vibrational frequencies. For a
non-linear molecule with n atoms, the number of fundamental vibrational modes is (3n-6); linear

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molecules have 3n-5 fundamental vibrational modes. Therefore, water - a non-linear molecule
theoretically possesses 3 fudamental vibrations – two stretching and one bending (figure 16); whereas
carbon dioxide - a linear molecule possess 4 fundamental absorption bands involving two stretching and
two bending modes (figure 17). Amongst these theoretically possible vibrations, a stretching vibration is
a rhythmical movement along the bond axis such that interatomic distance is increasing or decreasing. A
bending vibration consists of a change in bond angle between bonds with a common atom or the
movement of a group of atoms with respect to remaining part of the molecule without movement of the
atoms in the group with respect to one another.

Figure 16: Vibrational modes for water molecule

Figure 17: Vibrational modes for carbon dioxide molecule

The various stretching and bending modes can be represented by considering an AX 2 group
appearing as a portion of molecule, for example, the CH 2 group in a hydrocarbon molecule (figure 18 ).
Any atom joined to two other atoms will undergo comparable vibrations for example NH 2 or NO2. Each of
different vibration modes may give rise to a different absorption band so that CH2 groups give rise to two

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C-H stretching bands i.e. υsym and υantisym. Some of the vibrations may have the same frequency i.e.
they are degenerate and their absorption bands will appear at same position (for CO2, see figure 17).

Figure 18: Vibrational modes of a CH2 group. [ and indicate movement above and below the
plane of page]

In addition to the fundamental vibrations, other frequencies can be generated by modulations of

the fundamental bands. Overtone bands appear at integral multiples of fundamental vibrations.
Therefore, the strong absorptions at say 800 cm-1 and 1750 cm-1 will also give rise to weaker absorptions
at 1600 cm-1 and 3500 cm-1, respectively. In the IR spectra of benzaldehyde (figure 33) and acetophenone
(figure 31), due to C=O stretching vibration a weak overtone can be seen at 3400 and 3365 cm-1,

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respectively. Two frequencies may interact to give beats which are combination or difference frequencies.
The absorptions at x cm-1 and y cm-1 interact to produce two weaker beat frequencies at x + y cm-1 and x
– y cm-1. Therefore, whereas the factors like degeneracy of bands from several absorptions of the same
frequency, lack of change in molecular dipole moment during vibration and fall of frequencies outside the
4000-400 cm-1 region decrease the number of bands whereas the overtone and beats increase the number
of bands actually appeared in IR spectrum. Therefore, theoretical numbers of fundamental frequencies
are seldom observed. Other Factors Influencing Vibrational Frequencies The vibrational frequency of a
bond, being part of a molecule, is significantly affected by the electronic and steric factors of the
surroundings, in addition to the bond strength and atomic masses discussed above. When two bond
oscillators share a common atom, they seldom behave as individual oscillators where the individual
oscillation frequencies are widely different. The mechanical coupling interactions between two oscillators
are responsible for these changes.

For example, the carbon dioxide molecule, which consists of two C=O bonds with a common
carbon atom, has two fundamental stretching vibrations – an asymmetrical and a symmetrical stretching
mode. The symmetrical stretching mode produces no change in dipole moment and is IR inactive.
Asymmetric stretching mode is IR active and appears at a higher frequency (2350 cm-1) than observed for
a carbonyl group in aliphatic ketones (1715 cm-1). The carbonyl stretching frequency in RCOCH3 (~1720
cm-1) is lower than acid chloride RCOCl (1750-1820 cm-1). This change in frequency of the C=O stretching
may be arising due to (i) difference in mass between CH3 and Cl (ii) the inductive or mesomeric influence
of Cl on the C=O bond (iii) coupling interactions between C=O and C-Cl bonds (iv) change in bond angles
arising due to steric factors etc. It is usually impossible to isolate one effect from the other. However, the
appropriate emphasis can be placed on those features that seem to be most responsible in explaining the
characteristic appearance and position of group frequencies.

Sample Preparation

For recording an IR spectrum, the sample may be gas, a liquid, a solid or a solution of any of these.
The samples should be perfectly free of moisture, since cell materials (NaCl, KBr, CsBr etc.) are usually
spoiled by the moisture. Liquids are studied neat or in solution. In case of neat liquid, a thin film of < 0.01
mm thickness is obtained by pressing the liquid between two sodium chloride plates and plates are
subjected to IR beam. Spectra of solutions are obtained by taking 1-10 % solution of the sample in an
appropriate solvent in cells of 0.1-1 mm thickness. A compensating cell, containing pure solvent is placed
in the reference beam of the instrument. The choice of solvent depends on the solubility of the sample
and its own minimal absorption in IR region. Carbon tetrachloride, chloroform and carbon disulfide are
preferred solvents. The spectrum of a solid can be obtained either as a mull or as an alkali halide pellet.
Mulls are obtained by thoroughly grinding 2-5 mg of a solid sample with a drop of mulling agent usually
Nujol (mixture of parafinic hydrocarbons) or fluorolube (a completely fluorinate polymer). The suspended
particles must be less than 2 µM to avoid excessive scattering of radiations. The mull is placed between
two sodium chloride plates and plates are subjected to IR beam. For preparing, an alkali halide pellet, 1-2

25
mg of dry sample is grinded with ~ 100 mg of KBr powder. The mixture is then pressed into a transparent
pellet with a special die under a pressure of 10,000-15,000 psi. KBr pellet is then mounted on holder and
is placed in sample beam of IR spectrophotometer.

Characteristic Group Vibrations of Organic Molecules

An infrared spectrum of an organic compound comprises many bands and assigning each band to
a particular mode of vibration is practically impossible but two non-identical molecules generally have
different IR spectra. An IR spectrum, therefore, is a fingerprint of the molecule. The region most useful for
the purpose of “fingerprinting” of the compound is 650-1350 cm-1. This region comprises a large number
of bands due to skeletal vibrations and when the spectrum we are studying coincides exactly with the
spectrum of a known compound, it can be safely assumed that the two compounds are identical. The
region above 1350 cm-1 often provides easily recognizable bands of various functional groups and thus
much valuable structural evidence from relatively few of theses bands is obtained and total interpretation
of the complete spectrum is seldom required. In the following sections, the basic information about the
vibrational modes in basic functional groups has been discussed.

1. Hydrocarbons C-H and C-C stretching and bending vibrations

(i) Alkanes: In simple hydrocarbons, only two types of atoms - C and H and only two types of
bonds – C-C and C-H are present. The C-H starching vibrations usually occur in the general region between
3300 cm-1 (in alkynes) and 2700 cm-1 (in aldehydes).

A hydrocarbon containing a methyl group usually shows two distinct bands, one at 2960 cm-1 due
to asymmetric stretching and the other at 2870 cm-1 due to symmetric stretching. The C-H bonds in
methylene group undergo number of stretching and bending vibrations as shown in figure 18. The two
stretching vibrations – asymmetrical and symmetrical occur at 2925 cm-1 and appear in the spectrum
within a range of ± 10 cm-1.

The C-H bending vibrations of the methyl groups in the hydrocarbons normally occur at 1450 and
1375 cm-1. The band at 1375 cm-1 is due to methyl on the carbon atom and is quite sensitive to the
electronegativity of the substituent present at the methyl group. It shifts from as high as 1475 cm-1 in CH3-
F to as low as 1150 cm-1 in CH3-Br. However, this band is extremely useful in detecting the presence of
methyl group in a compound because it is sharp and of medium intensity and is rarely overlapped by
absorptions due to methylene or methine deformations. The intensity of this band usually increases with
the number of methyl groups in the compound. However, the presence of two or more methyl groups on
one aliphatic carbon atom (isopropyl or t-butyl groups) results in splitting of this band due to in-phase or
out-of phase interactions of the two symmetrical methyl deformations.

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 In case of methylene group, C-H bending vibrations such as scissoring, twisting, wagging and
rocking normally appear at fairly different frequencies. If two or more CH2 groups are present, the usually
strong scissoring and rocking absorption bands appear at 1465 and 720 cm-1, respectively. Whereas weak
bands due to twisting and wagging vibrations appear at 1250 + 100 cm -1. So, the scissoring absorption
band of methylene around 1465 cm-1 often overlaps with asymmetric bending vibration of methyl at 1450
cm-1.

In cyclic aliphatic hydrocarbons, the C-H stretching frequencies are the same (2800 – 3000 cm-1)
as in the case of acyclic compounds, if the ring is unstrained. However, methylene scissoring bands shift
slightly to smaller wavenumber (1470 cm-1 in hexane and 1448 cm-1 in cyclohexane, see figure 19). In
sterically strained cyclic compounds, the C-H stretching normally occurs at slightly higher wavenumber
e.g. 3080 -3040 cm-1 in cyclopropane.

The C-C bond vibrations appear as weak bands in 1200-800 cm-1 region and are seldom used for
structural study. Whereas the C-C bending absorptions occur at < 500 cm-1 and are usually below the range
of IR – instrument.

Figure 19: The infrared spectrum of cyclohexane (neat liquid)

(ii) Alkenes: The carbon-carbon double bond has a higher force constant than a C-C single bond
and in a non-conjugated olefin, C=C stretching vibrations appear at higher frequency (1680-1620 cm-1 )
than that of the C-C starching vibrations (1200-800 cm-1).

In completely symmetrical alkenes, such as ethylene, tetrachloroethylene etc., C=C stretching

band is absent, due to lack of change in dipole moment in completely symmetrical molecule. On the other
hand, non-symmetrically substituted double bonds exhibit strong absorption bands. The absorption bands
are more intense for cis isomers than for trans isomers; for mono or tri substituted olefins than for di and
tetra substituted ones. Also, terminal olefins show stronger C=C double bond stretching vibrations than
internal double bonds. Similarly C=C groups conjugated with certain unsaturated group show stronger

27
band than for non-conjugated ones. In case of olefins, conjugated with an aromatic ring, the C=C
stretching appears at 1625 cm-1 (s) and an additional band at ~1600 cm-1 is observed due to aromatic
double bond. In compounds containing both olefinic and alkyl C-H bonds, the bands above 3000 cm-1 are
generally attributed to aromatic or aliphatic C-H stretching, whereas between 3000-2840 cm-1 are
generally assigned to the alkyl C-H stretching. The absorption frequency of a double bond in a cyclic ring
is very sensitive to ring size (figure 20). The absorption frequency decreases as the internal angle
decreases and is lowest in cyclobutene (90o angle). The frequency increases again for cyclopropane.

Figure 20: C=C vibration frequencies of cycloalkenes

The exocyclic double bonds exhibit an increase in frequency with decrease in ring size (figure 21). The
exocyclic double bond on six-membered ring absorbs at 1651 cm-1 and it is shifted to 1780 cm-1 in case of
exocyclic double bond on cyclopropane. The allenes show the highest double bond absorptions at 1940
cm-1.

Figure 21: Exocyclic C=C double bond frequencies in various ring sizes

(iii) Alkynes : All alkynes both terminal (R.C≡CH ) or non-terminal (R.C≡CR ) contain carbon – carbon triple
bond but the non-terminal alkynes also contain a ≡CH bond. The force constant for a triple bond is grater
than that for a double bond. Consequently, whereas a C-C stretching vibrations occur between 1300-800
cm-1 and the C=C stretching vibration occur in the region 1700-1500 cm-1, the C≡C vibrations are observed
at significantly higher frequencies in the region of 2300 to 2050 cm-1. The terminal alkynes show weak
triple bond stretching vibration at 2140-2050 cm-1, whereas the unsymmetrically disubstituted alkynes
show a triple bond absorption at 2260-2190 cm-1. The acetylene C-H stretching vibrations are normally
observed as a sharp characteristic band in the region 3310-3200 cm-1 and the acetylenic C-H bending

28
vibrations occur in the region 600-650 cm-1. Therefore the frequency of the absorption of C-H bond is a
function of the type of hybridization of the carbon to which hydrogen atom is attached. While moving
from sp3 to sp2 and sp hybridized carbons, the s-character increases and so is the bond strength (force
constant of C-H bond and the frequency of absorption (Table 7).

Table 7:

(iv) Aromatic Hydrocarbons: In the aromatic compounds, the most prominent bands are due to out-of-
plane bending of ring C-H bonds in the region of 900-650 cm-1. These bands can be used to assign the ring
substitution pattern in mono substituted benzenes and 1,2-, 1,3-, and 1,4- substituted benzene
derivatives. Mono substituted benzene derivatives exhibit strong absorption band near 690 cm-1 (see IR
spectrum of toluene, figure 22). The absence of this band indicates the absence of mono substituted
phenyl rings in the sample. A second strong band appears at ~750 cm-1. 1,2-Disubstituted benzenes give
one strong absorption band at ~ 750 cm-1. 1,3- Disubstituted rings give three absorption bands at ~690,
~780 and ~ 880 cm-1. 1,4-Disubstituted rings give one strong absorption band in the region 800-850 cm-1
(strong absorption band at 831 cm-1 is seen in IR spectrum of t-butylphenol, figure 26 ). The spectra of

29
aromatic compounds typically exhibit many weak or medium intensity C-H stretching vibrations in the
region 3100- 3030 cm-1, the region of olefinic compounds.

Figure 22: The infrared spectrum of toluene (neat liquid)

Figure 23: The infrared spectrum of mesitylene (neat liquid)

The bands considered to be of most help in diagnosing the aromatic character of the compound appear
in the region 1650-1400 cm-1. There are normally four bands in this region at about 1600, 1585, 1500 and
1450 cm-1 and are due to C=C in-plane vibrations (see spectra in figures 22 and 23). The combination and
overtone bands in 2000-1650 cm-1 region are also characteristics of aromatic rings. Moreover, they are
very weak and are observed only in the case of concentrated solutions of highly symmetric benzene
derivatives.

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2. Alcohols and Phenols

When a hydrogen atom from an aliphatic hydrocarbon is replaced by an OH group, new bands
corresponding to new OH and C-O band absorption appear in the IR spectrum. A medium to strong
absorption band from 3700 to 3400 cm-1 ( see IR spectra of 1-butanol and t-butylphenol in figures 25 and
26) is a strong indication that the sample is an alcohol or phenol (The presence of NH or moisture causes
similar results). The exact position and shape of this band depends largely on the degree of H-bonding. A
strong, sharp peak in the region as higher 3700 cm-1 in gaseous or extremely dilute solutions represents
unbounded or free OH group(s). Alcohols and phenols in condensed phases (bulk liquid, KBr discs,
concentrated solution etc.) are strongly hydrogen bonded, usually in the form of dynamic polymeric
association; dimmers, trimers, tetramers etc. (Figure 24) and cause broadened bands at lower
frequencies. The hydrogen bonding involves a lengthening of the original O-H bond. This bond is
consequently weakened, force constant is reduced and so the stretching frequency is lowered.

Figure 24: Polymeric association of ROH

The C-O stretching in phenols / alcohols occurs at a lower frequency range 1250-1000 cm-1. The
coupling of C-O absorption with adjacent C-C stretching mode, makes it possible to differentiates between
primary (~1050 cm-1), secondary (~1100 cm-1) and tertiary ( ~1150 cm-1) alcohols and phenols ( ~1220 cm-
1).

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 Figure 25: The infrared spectrum of 1-butanol (neat liquid)

Figure 26: The infrared spectrum of t-butylphenol (nujol mull)

3. Carbonyl Compounds The absorption peak for C=O stretching in the region 1870 to 1600 cm-1
is perhaps the easiest band to recognize in IR spectrum and is extremely useful in analysis of carbonyl
compounds. The changes in C=O stretching frequency in various carbonyl compounds viz. aldehydes,
ketones, acids, esters, amides, acid halides, anhydrides etc. can be explained by considering (i) electronic
and mass effects of neighboring substituents (ii) resonance effects ( both C=C and heteroatom lone pair)
(iii) hydrogen bonding (inter and intramolecular) (iv) ring strain etc. It is customary to refer to the
absorption frequency of a saturated aliphatic ketone at 1715 cm-1 as normal value and changes in the

32
environment of the carbonyl group can either lower or raise the absorption frequency from the “normal”
value.

(i) Inductive and Resonance Effects: The replacement of an alkyl group of the saturated aliphatic
ketone by a heteroatom (O, N) shifts the C=O stretching frequencies due to inductive and resonance
effects. In esters, the oxygen due to inductive effect withdraws the electrons from carbonyl group (figure
27) and increases the C=O bond strength and thus the frequency of absorption. In amides, due to the
conjugation of lone pair of electrons on nitrogen atom, the resonance effect increases the C=O bond
length and reduces the C=O absorption frequency. Therefore, C=O absorption frequencies due to
resonance effects in amides are lowered but due to inductive effect in esters are increased than those
observed in ketones.

Figure 27: Inductive and resonance effects in ester and amide groups

In acid chlorides, the halogen atom strengthens the C=O bond through inductive effect and shifts
the C=O stretching frequencies even higher than are found in esters. The acid anhydrides give two bands
in C=O stretching frequency region due to symmetric (~1820 cm-1) and asymmetric (~1760 cm-1)
stretching vibrations (figure 28).

Figure 28: Inductive effect in acid chloride and C=O stretch in anhydride

(ii) Conjugation Effects: The C=O stretching frequencies for carbon-carbon double bond conjugated
systems are generally lower by 25-45 cm-1 than those of corresponding nonconjugated compounds. The
delocalization of π-electrons in the C=O and C=C bonds lead to partial double bond character in C=O and
C=C bonds and lowers the force constant (figure 30). Greater is the ability of delocalization of electrons,
the more is lowering in C=O stretching frequency. In general s-cis conformations absorb at higher

33
frequency than s-trans conformations. A similar lowering in C=O stretching frequency occurs when an aryl
ring is conjugated with carbonyl compound.

Figure 30: Resonance effects and s-cis and s-trans structures in enones

(iii) Ring Size Effects: Six-membered rings with carbonyl group e.g. cyclohexanone absorb at normal value
i.e. 1715 cm-1. Decrease in ring size increases the C=O stretching frequency. Smaller rings require the use
of more p- character to make C-C bonds for the requisite small angles. This gives more s character to the
C=O sigma bond and thus results in strengthening of C=O double bond. The comparison of C=O stretching
frequencies of various compounds in figure 31 shows that in ketones and esters, ~ 30 cm-1 increase in
frequency occurs on moving to one carbon lower ring.

Figure 31: C=O stretching frequencies in various compounds

(iv) Hydrogen Bonding Effects: Hydrogen bonding to a C=O group withdraws electrons from oxygen and
lowers the C=O double bond character. This results in lowering of C=O absorption frequency. More
effective is the hydrogen bonding, higher will be the lowering in C=O absorption frequencies.

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Figure 32: H-bonding effects on C=O stretch

The monomeric carboxylic acids (in very dilute solutions) absorb at ~1760 cm -1. The dimerization of
carboxylic acids in their concentrated solutions or in solid state lowers the carboxyl carbonyl frequency to
1710 cm-1. The more effective intramolecular hydrogen bonding in methyl salicylate lowers the C=O
stretching frequency to 1680 cm-1 than observed at 1700 cm-1 in case of methyl p-hydroxybenzoate.

a. Aldehydes and Ketones

Aliphatic aldehydes show strong C=O stretching in the region of 1740 – 1725 cm-1. The conjugation
of an aldehyde to a C=C or a phenyl group lowers C=O stretching by ~ 30 cm -1. This effect is seen in
benzaldehyde in which aryl group is attached directly to the carbonyl group and shifts C=O stretch to
1701 cm-1 (see IR spectrum of benzaldehyde, figure 33 ). Aldehyde C-H stretching vibrations appear
as a pair of weak bands between 2860-2800 and 2760-2700 cm-1. The higher C-H stretching band
(2860-2800 cm-1) of aldehyde is often buried under aliphatic CH band. But the lower C-H band at 2760-
2700 cm-1 is usually used to distinguish aldehydes from ketones. The C-H bending vibrations appear
between 945-780 cm-1.

Figure 33: The infrared spectrum of benzaldehyde (neat liquid)

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The aliphatic acyclic ketones show C=O stretching between 1720 to 1700 cm-1 which is shifted to lower
frequencies by 20-30 cm-1 on conjugation with C=C or phenyl ring. The presence of two conjugated groups
as in benzophenone further lowers the C=O stretching frequency to 1665 cm-1 (figures 34 and 35 ).

Figure 34: The infrared spectrum of acetophenone (neat liquid)

Figure 35: The infrared spectrum of benzophenone (nujol mull)

36
In case of cyclic ketones, the coupling between C=O stretching and C(=O)-C single bond causes increase in
C=O stretching frequency as the C-C(=O) angle decreases (figure 31).

b. Carboxylic Acids, Esters and Carboxylates In case of carboxylic acids, in solid state or pure liquid state,
the intermolecular hydrogen bonding weakens the C=O bond and thus lower the stretching frequency to
~1720 cm-1. The O-H stretch appears as a very broad band between 3400 – 2500 cm-1 (see IR spectrum of
benzoic acid, figure 36). The appearance of strong C=O stretching along with broad hydroxyl peak
centered at ~ 3000 cm-1 in an IR spectrum certainly shows the presence of carboxylic acid. In addition a
medium intensity C=O stretch appears between 1320 – 1260 cm-1. In dilute solutions, the carboxylic acids
attain monomeric structures and the inductive effect of oxygen shifts the C=O absorption band to higher
values (1760 – 1730 cm-1) than observed in ketones.

Figure 36: The infrared spectrum of benzoic acid (nujol mull)

In case of esters, the C=O stretching appears in the range 1750-1730 cm-1 and strong C-O stretching
absorption appears in the range 1300 -1000 cm-1. The esters of α,β-unsaturated or aryl carboxylic acids
due to conjugation absorb at lower frequency (figure 37).

Figure 37: Effect of conjugation in cyclic esters

ROTATIONAL SPECTROSCOPY (MICROWAVE SPECTROSCOPY)

37
Introduction

Free atoms do not rotate or vibrate. For an oscillatory or a rotational motion of a pendulum, one
end has to be tied or fixed to some point. In molecules such a fixed point is the center of mass. The atoms
in a molecule are held together by chemical bonds. The rotational and vibrational energies are usually
much smaller than the energies required to break chemical bonds. The rotational energies correspond to
the microwave region of electromagnetic radiation (3x1010 to 3x1012 Hz; energy range around 10 to 100
J/mol) and the vibrational energies are in the infrared region (3x1012 to 3x1014 Hz; energy range around
10kJ/mol) of the electromagnetic radiation. For rigid rotors (no vibration during rotation) and harmonic
oscillators (wherein there are equal displacements of atoms on either side of the center of mass) there
are simple formulae characterizing the molecular energy levels. In real life, molecules rotate and vibrate
simultaneously and high speed rotations affect vibrations and vice versa. However, in our introductory
view of spectroscopy we will simplify the picture as much as possible. We will first take up rotational
spectroscopy of diatomic molecules.

Rotational Spectra of diatomics

Figure 38: A rigid diatomic with masses m1 and m2 joined by a thin rod of length r = r 1 + r2 .The centre of
mass is at C.

The two independent rotations of this molecule (Figure 38) are with respect to the two axes which
pass though C and are perpendicular to the “bond length” r. The rotation with respect to the bond axis is
possible only for “classical” objects with large masses. For quantum objects, a “rotation” with respect to
the molecular axis does not correspond to any change in the molecule as the new configuration is
indistinguishable from the old one.

The center of mass is defined by equating the moments on both segments of the molecular axis.

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The quantized rotational energy levels for this diatomic are

The energy differences between two rotational levels is usually expressed in cm-1. The wave number
corresponding to a given ∆E is given by

The rotational energy levels of a diatomic molecule are shown in Figure 39.

39
Figure 39: Rotational energy levels of a rigid diatomic molecule and the allowed transitions.

The selection rule for a rotational transition is,

In addition to this requirement, the molecule has to possess a dipole moment. As a dipolar molecule
rotates, the rotating dipole constitutes the transition dipole operator μ. Molecules such as HCl and CO will
show rotational spectra while H2, Cl2 and CO2 will not. The rotational spectrum will appear as follows
(Figure 40).

Figure 40: Rotational spectrum of a rigid diatomic. Values of B are in cm-1. Typical values of B in cm-1
are 1.92118 (CO), 10.593 (HCl), 20.956 (HF)

By using rotational or microwave spectroscopy, very accurate values of bond lengths can be
obtained. For example, in HCN, the C-H length is 0.106317 ± 0.000005 nm and the CN bond length is
0.115535 ± 0.000006 nm. The principle of the microwave oven involves heating the molecules of water

40
 through high speed rotations induced by microwaves. The glass container containing water
however remains cold since it does not contain rotating dipoles.

41
 Question Bank for Module IV

1. What is absorption spectroscopy? Explain with suitable examples.

2. What is emission spectroscopy? Explain with example.
3. Discuss the terms ‘absorbance’ and ‘transmittance’ of monochromatic light through an
absorbing medium.
4. What is flame emission spectroscopy? What are its disadvantages?
5. What is ICP-AES? Briefly explain its working.
6. What is Lambert-Beer’s law? Explain.
7. How Lambert-Beer’s law is the fundamental backbone behind quantitative estimation of an
analyte?
8. What is molar extinction coefficient?
9. What are the limitations of Lambert-Beer’s law?
10. With the help of a neat diagram, compare the electronic transitions possible for C-C, C=C, C=O
& C=C-C=C bonds.
11. Discuss ‘solvent effect’ in case of UV-VIS spectroscopy of organic molecules.
12. What is a chromaphore? Give examples of a few simple unconjugated chromaphores and their
electronic transitions.
13. What is an auxochrome? Give examples.
14. Discuss the effect of presence of conjugated double bond on the MO of an organic molecule.
15. Discuss the concept of infrared spectroscopy.
16. What do you mean by ‘fingerprint region’ of IR absorption spectrum of an organic molecule?
17. Discuss the concept of ‘reduced mass’ in case of a diatomic molecule.
18. Derive the mathematical expression for ‘wave number’ of vibration of a polar bond.
19. Which of the following compounds can be estimated by IR spectroscopy? (i) N 2 (ii) CO (iii) CH3OH
(iv) O2 (v) H2O
20. Draw the vibration modes for water molecule.
21. Draw the vibration modes for carbon monoxide molecule.
22. Discuss the sample preparation methods in IR spectroscopy.
23. Why KBr is used as matrix in sample preparation for IR spectroscopy?
24. Discuss the various stretching and bending vibrations in the following group of compounds and
their approximate wave number: (i) alkanes (ii) alkenes (iii) alkynes (iv) aromatic hydrocarbons
(v) alcohols and phenols
25. Discuss ‘ring size effect’ in carbonyl compounds.
26. What is the effect of hydrogen bonding on the IR spectrum of a carbonyl compound?
27. Write the expression for quantized rotational energy levels of a diatomic molecule.
28. Discuss the concept of proton NMR.
29. What do you mean by the term ‘chemical shrift’?
30. Discuss the proton NMR signals for the following: (i) CH3CH2CH2 Br (ii) C2H5OH.

42