Final year Bpharm 7th semester
Vikrant Ashok Shirke
Gsmcop, Wagholi, Pune
BP701T- INSTRUMENTAL METHODS OF ANALYSIS.
UNIT-I (28-44 MARKS)
Syllabus:-
UV-Visible Spectroscopy:- Introduction to spectroscopy, electronic transitions, chromophores, auxochromes,
spectral shifts, solvent effects on absorption spectra, Beer and Lambert’s law, derivation and deviations.
Instrumentation:- Sources of radiation, wavelength selectors, sample cells.
Detectors:- Photo tube, photomultiplier tube, photo voltaic cell, Silicon photodiode.
Application:- Spectrophotometric titrations, single components and multi components analysis.
Fluorimetry:- Theory, concept of singlet, doublet, and triplet electronic states, internal and external conversions,
factors affecting fluorescence, quenching, instrumentation and applications.
Q.1. Attempt the following (3 Marks Each)
1) Discuss in detail the various types of transitions involved in UV-Visible Spectroscopy? (3/5Marks)
➢ In UV-Visible spectroscopy, transitions refer to the movement of electrons from one
energy level to another when they absorb ultraviolet (UV) or visible (Vis) light, therefore
different types of transitions are observed based on the nature of the molecule and the
type of electrons involved.
1) σ → σ* (Sigma to Sigma-Antibonding Transitions):-
Electrons in a σ bonding orbital moving to an σ* antibonding orbital.
They are Found in saturated hydrocarbons and compounds with single bonds.
They have very high energy (short wavelength, ~125–150 nm).
Not observed in normal UV spectra due to very high energy requirement.
Least important for UV-Vis spectroscopy as it lies outside the accessible range.
[ Example: Methane (CH₄), ethane (C₂H₆) ]
2) n → σ* (Non-bonding to Sigma-Antibonding Transition):-
A non-bonding electron pair (n) moving to a σ* antibonding orbital.
They Seen in molecules with lone pairs.
They required Less than σ → σ*, occurs in the near UV region (~150–250 nm).
Common in saturated compounds with heteroatoms (O, N, halogens).
[ Example: Water (H₂O), ammonia (NH₃) ]
3) π → π* (Pi to Pi-Antibonding Transition):-
A π bonding electron (from double or triple bonds)moving to a π* antibonding orbital.
It is Common in aromatic systems, and compounds with C=C, C≡C, or C=O.
It requires moderate energy, typically occurs in the (200–400 nm range).
Most common transition in conjugated organic systems.
[ Example: Ethylene, benzene ]
4) n → π* (Non-bonding to Pi-Antibonding Transition):-
A lone pair (n) electron being excited to a π* antibonding orbital.
It is Common in carbonyl compounds, nitro compounds.
, They require Lower than π → π*, usually occurs around (250–400nm range)
Important in detecting carbonyl-containing compounds in UV range.
[ Example: Formaldehyde, acetone. ]
2) Explain with example the excitation and emission fluorescence spectra?
➢ Fluorescence is a type of luminescence where a molecule absorbs light at one
wavelength (excitation) and emits light at a longer wavelength (emission). This
phenomenon is commonly studied using fluorescence spectroscopy.
1) Excitation Spectrum:-
The excitation spectrum shows the intensity of fluorescence as a function of the excitation
wavelength, while keeping the emission wavelength constant.
It helps identify which excitation wavelengths are most efficient in producing fluorescence.
➤ Procedure:
1)The emission monochromator is fixed at a specific emission wavelength (λem).
2)The excitation monochromator is scanned over a range of excitation wavelengths (λex).
3) The resulting spectrum shows how efficiently different wavelengths excited the sample.
➤ Appearance: It resembles the absorption spectrum, because fluorescence occurs only when
the molecule absorbs light.
2) Emission Spectrum:-
The emission spectrum shows the intensity of fluorescence as a function of the emission
wavelength, while keeping the excitation wavelength constant.
It reveals the wavelengths of light emitted by the excited molecule.
➤ Procedure:
1)The excitation wavelength (λex) is fixed.
2) The emission monochromator scans over various emission wavelengths (λem).
3) The spectrum displays the characteristic fluorescence emission.
➤ Appearance: This spectrum typically shows a broad peak at longer wavelengths (lower energy)
than the excitation wavelength due to Stokes shift.
[ Example: Fluorescein ]
Fluorescein is a well-known fluorescent dye used in biochemical applications.
➤ Excitation Spectrum:-
● Fixed Emission: λem= 520 nm • Scan Excitation: 400–500 nm
● Peak Excitation Wavelength: ~495 nm
➤ Emission Spectrum:-
● Fixed Excitation: λex= 495 nm •Scan Emission: 500–600 nm
● Peak Emission Wavelength: ~520 nm
This shows a Stokes shift of about 25 nm (difference between excitation and emission maxima).
3) Discuss the effects of solvent on absorption spectra in UV spectroscopy? (3/5Marks)
➢ In UV-Visible spectroscopy, the solvent used can significantly affect the absorption
spectra of a compound. These effects are known as solvent effects, and they primarily
influence the position (wavelength) and intensity (absorbance) of absorption bands.
1) Solvent Polarity and Electronic Transitions
a) π → π* Transitions
In polar solvents, excited states (π*) are more stabilized than the ground states.
This leads to a bathochromic shift (red shift or shift to longer wavelength).
,[ Example: Ethylene in hexane vs. ethanol — the π → π* band shifts to longer wavelengths in
ethanol. ]
b) n → π* Transitions
Lone pair electrons (n) interact with polar solvents via hydrogen bonding or dipole interactions.
This interaction stabilizes the ground state more than the excited state.
This leads to hypsochromic shift (blue shift or shift to shorter wavelength).
[Example: Acetone in water vs. hexane — the n → π* band shifts to shorter wavelengths in water.
]
2) Solvent Effects on Band Intensity
π → π* transitions often show increased intensity in polar solvents.
n → π* transitions may become weaker in hydrogen-bonding solvents due to partial removal of
non-bonding electrons (less availability for transition).
3) Solvent–Solute Interactions
These include Hydrogen bonding, Dipole-dipole interactions, Van der Waals forces,such
interactions can lead to broadening of bands or changes in spectral shape.
4) Practical Implications in Spectroscopy
Choose solvents that do not absorb in the same region as the analyte.
Use nonpolar solvents (like hexane) when minimizing solvent effects is desired.
Use polar solvents to study solvent interactions or enhance specific transitions.
5) Hydrogen Bonding
Solvents capable of hydrogen bonding (e.g, water, alcohols) can interact with functional groups
like –OH, –NH₂, or C=O in the solute.
These interactions may lead to broadening of the absorption band or slight shifts in λmax.
6) Solvent Transparency
The solvent must not absorb in the same UV region as the analyte.
● Common UV-transparent solvents:
a) Non-polar: Hexane, cyclohexane
b) Polar: Ethanol, methanol, acetonitrile
c) High UV cutoff solvents (to be avoided): Water, chloroform.
7) Solvent Effects on Band Intensity
Solvent polarity can also affect molar absorptivity (ε) by changing the transition dipole moment,
leading to changes in absorption intensity.
Hence, Solvent choice in UV spectroscopy is critical, it affects Wavelength position (λmax),
Absorption intensity, Band shape and width.
Therefore, Understanding solvent effects helps in interpreting spectra correctly and designing
accurate experiments.
4) Difference between fluorescence and phosphorescence phenomenon?
Feature Fluorescence Phosphorescence
Type of Emission Immediate light emission Delayed light emission
Time Duration Lasts for nanoseconds (10⁻⁹ to Lasts from milliseconds to
10⁻⁸ sec) minutes or even hours
Excited State Transition from singlet excited Transition from triplet excited
state (S₁ → S₀) state (T₁ → S₀)
, Transition Type Spin-allowed (fast) Spin-forbidden (slow)
After Removal of Stops almost immediately Continues glowing after
Excitation removal of light source
Energy of Emitted Higher energy (shorter Lower energy (longer
Light wavelength) wavelength)
Example Fluorescent dyes, vitamins (like Glow-in-the-dark materials
riboflavin) (e.g., strontium aluminate)
Visual Analogy the drop is "straight down" the electron detours through a
( Jablonski Diagram) “side trip” to a triplet state.
Type of Transition Singlet → Singlet. Triplet → Singlet.
5) Explain the term chromophore and auxochrome with suitable example? (3/5Marks)
Auxochrome–Chromophore Theory:-
➢ The Auxochrome-Chromophore Theory is fundamental in understanding the absorption
of ultraviolet (UV) and visible light by organic molecules. This theory helps explain why
certain compounds absorb light at specific wavelengths, and how structural modifications
influence their color and spectral properties.
1) Chromophore:- A chromophore is a part of a molecule responsible for absorbing UV or
visible light.
It typically contains π-electrons or non-bonding electrons (n) that can undergo electronic
transitions such as π → π*, or n → π*.
Absorption of light by chromophores promotes electrons from the ground state to an excited
state, resulting in an absorption band in the UV-visible spectrum.
● Examples of Chromophores:
Chromophore Common Structure Transition Type
C=C (Alkene) –CH=CH– π → π*
C=O (Carbonyl) –C=O n π*, π → π*
NO₂ (Nitro group) –NO₂ n π*, π → π*
Aromatic Rings Benzene-like structures π → π*
2) Auxochrome:- An auxochrome is a group that, when attached to a chromophore, does
not absorb light by itself in the visible region, but modifies the ability of the chromophore
to absorb light.
It usually contains lone pair electrons (e.g, –OH, –NH₂, –OR, –Cl).
● Auxochromes affect absorption in two main ways:
a) Shifting λmax (wavelength of maximum absorption)
b) Increasing intensity of absorption (ε value)
● Types of Shifts:
Shift Type Description. Caused by
Bathochromic shift. Absorption moves Electron-donating auxochromes
to longer
wavelength (red)
Hypsochromic shift. Absorption moves Electron-withdrawing auxochromes
to shorter
wavelength (blue)
Vikrant Ashok Shirke
Gsmcop, Wagholi, Pune
BP701T- INSTRUMENTAL METHODS OF ANALYSIS.
UNIT-I (28-44 MARKS)
Syllabus:-
UV-Visible Spectroscopy:- Introduction to spectroscopy, electronic transitions, chromophores, auxochromes,
spectral shifts, solvent effects on absorption spectra, Beer and Lambert’s law, derivation and deviations.
Instrumentation:- Sources of radiation, wavelength selectors, sample cells.
Detectors:- Photo tube, photomultiplier tube, photo voltaic cell, Silicon photodiode.
Application:- Spectrophotometric titrations, single components and multi components analysis.
Fluorimetry:- Theory, concept of singlet, doublet, and triplet electronic states, internal and external conversions,
factors affecting fluorescence, quenching, instrumentation and applications.
Q.1. Attempt the following (3 Marks Each)
1) Discuss in detail the various types of transitions involved in UV-Visible Spectroscopy? (3/5Marks)
➢ In UV-Visible spectroscopy, transitions refer to the movement of electrons from one
energy level to another when they absorb ultraviolet (UV) or visible (Vis) light, therefore
different types of transitions are observed based on the nature of the molecule and the
type of electrons involved.
1) σ → σ* (Sigma to Sigma-Antibonding Transitions):-
Electrons in a σ bonding orbital moving to an σ* antibonding orbital.
They are Found in saturated hydrocarbons and compounds with single bonds.
They have very high energy (short wavelength, ~125–150 nm).
Not observed in normal UV spectra due to very high energy requirement.
Least important for UV-Vis spectroscopy as it lies outside the accessible range.
[ Example: Methane (CH₄), ethane (C₂H₆) ]
2) n → σ* (Non-bonding to Sigma-Antibonding Transition):-
A non-bonding electron pair (n) moving to a σ* antibonding orbital.
They Seen in molecules with lone pairs.
They required Less than σ → σ*, occurs in the near UV region (~150–250 nm).
Common in saturated compounds with heteroatoms (O, N, halogens).
[ Example: Water (H₂O), ammonia (NH₃) ]
3) π → π* (Pi to Pi-Antibonding Transition):-
A π bonding electron (from double or triple bonds)moving to a π* antibonding orbital.
It is Common in aromatic systems, and compounds with C=C, C≡C, or C=O.
It requires moderate energy, typically occurs in the (200–400 nm range).
Most common transition in conjugated organic systems.
[ Example: Ethylene, benzene ]
4) n → π* (Non-bonding to Pi-Antibonding Transition):-
A lone pair (n) electron being excited to a π* antibonding orbital.
It is Common in carbonyl compounds, nitro compounds.
, They require Lower than π → π*, usually occurs around (250–400nm range)
Important in detecting carbonyl-containing compounds in UV range.
[ Example: Formaldehyde, acetone. ]
2) Explain with example the excitation and emission fluorescence spectra?
➢ Fluorescence is a type of luminescence where a molecule absorbs light at one
wavelength (excitation) and emits light at a longer wavelength (emission). This
phenomenon is commonly studied using fluorescence spectroscopy.
1) Excitation Spectrum:-
The excitation spectrum shows the intensity of fluorescence as a function of the excitation
wavelength, while keeping the emission wavelength constant.
It helps identify which excitation wavelengths are most efficient in producing fluorescence.
➤ Procedure:
1)The emission monochromator is fixed at a specific emission wavelength (λem).
2)The excitation monochromator is scanned over a range of excitation wavelengths (λex).
3) The resulting spectrum shows how efficiently different wavelengths excited the sample.
➤ Appearance: It resembles the absorption spectrum, because fluorescence occurs only when
the molecule absorbs light.
2) Emission Spectrum:-
The emission spectrum shows the intensity of fluorescence as a function of the emission
wavelength, while keeping the excitation wavelength constant.
It reveals the wavelengths of light emitted by the excited molecule.
➤ Procedure:
1)The excitation wavelength (λex) is fixed.
2) The emission monochromator scans over various emission wavelengths (λem).
3) The spectrum displays the characteristic fluorescence emission.
➤ Appearance: This spectrum typically shows a broad peak at longer wavelengths (lower energy)
than the excitation wavelength due to Stokes shift.
[ Example: Fluorescein ]
Fluorescein is a well-known fluorescent dye used in biochemical applications.
➤ Excitation Spectrum:-
● Fixed Emission: λem= 520 nm • Scan Excitation: 400–500 nm
● Peak Excitation Wavelength: ~495 nm
➤ Emission Spectrum:-
● Fixed Excitation: λex= 495 nm •Scan Emission: 500–600 nm
● Peak Emission Wavelength: ~520 nm
This shows a Stokes shift of about 25 nm (difference between excitation and emission maxima).
3) Discuss the effects of solvent on absorption spectra in UV spectroscopy? (3/5Marks)
➢ In UV-Visible spectroscopy, the solvent used can significantly affect the absorption
spectra of a compound. These effects are known as solvent effects, and they primarily
influence the position (wavelength) and intensity (absorbance) of absorption bands.
1) Solvent Polarity and Electronic Transitions
a) π → π* Transitions
In polar solvents, excited states (π*) are more stabilized than the ground states.
This leads to a bathochromic shift (red shift or shift to longer wavelength).
,[ Example: Ethylene in hexane vs. ethanol — the π → π* band shifts to longer wavelengths in
ethanol. ]
b) n → π* Transitions
Lone pair electrons (n) interact with polar solvents via hydrogen bonding or dipole interactions.
This interaction stabilizes the ground state more than the excited state.
This leads to hypsochromic shift (blue shift or shift to shorter wavelength).
[Example: Acetone in water vs. hexane — the n → π* band shifts to shorter wavelengths in water.
]
2) Solvent Effects on Band Intensity
π → π* transitions often show increased intensity in polar solvents.
n → π* transitions may become weaker in hydrogen-bonding solvents due to partial removal of
non-bonding electrons (less availability for transition).
3) Solvent–Solute Interactions
These include Hydrogen bonding, Dipole-dipole interactions, Van der Waals forces,such
interactions can lead to broadening of bands or changes in spectral shape.
4) Practical Implications in Spectroscopy
Choose solvents that do not absorb in the same region as the analyte.
Use nonpolar solvents (like hexane) when minimizing solvent effects is desired.
Use polar solvents to study solvent interactions or enhance specific transitions.
5) Hydrogen Bonding
Solvents capable of hydrogen bonding (e.g, water, alcohols) can interact with functional groups
like –OH, –NH₂, or C=O in the solute.
These interactions may lead to broadening of the absorption band or slight shifts in λmax.
6) Solvent Transparency
The solvent must not absorb in the same UV region as the analyte.
● Common UV-transparent solvents:
a) Non-polar: Hexane, cyclohexane
b) Polar: Ethanol, methanol, acetonitrile
c) High UV cutoff solvents (to be avoided): Water, chloroform.
7) Solvent Effects on Band Intensity
Solvent polarity can also affect molar absorptivity (ε) by changing the transition dipole moment,
leading to changes in absorption intensity.
Hence, Solvent choice in UV spectroscopy is critical, it affects Wavelength position (λmax),
Absorption intensity, Band shape and width.
Therefore, Understanding solvent effects helps in interpreting spectra correctly and designing
accurate experiments.
4) Difference between fluorescence and phosphorescence phenomenon?
Feature Fluorescence Phosphorescence
Type of Emission Immediate light emission Delayed light emission
Time Duration Lasts for nanoseconds (10⁻⁹ to Lasts from milliseconds to
10⁻⁸ sec) minutes or even hours
Excited State Transition from singlet excited Transition from triplet excited
state (S₁ → S₀) state (T₁ → S₀)
, Transition Type Spin-allowed (fast) Spin-forbidden (slow)
After Removal of Stops almost immediately Continues glowing after
Excitation removal of light source
Energy of Emitted Higher energy (shorter Lower energy (longer
Light wavelength) wavelength)
Example Fluorescent dyes, vitamins (like Glow-in-the-dark materials
riboflavin) (e.g., strontium aluminate)
Visual Analogy the drop is "straight down" the electron detours through a
( Jablonski Diagram) “side trip” to a triplet state.
Type of Transition Singlet → Singlet. Triplet → Singlet.
5) Explain the term chromophore and auxochrome with suitable example? (3/5Marks)
Auxochrome–Chromophore Theory:-
➢ The Auxochrome-Chromophore Theory is fundamental in understanding the absorption
of ultraviolet (UV) and visible light by organic molecules. This theory helps explain why
certain compounds absorb light at specific wavelengths, and how structural modifications
influence their color and spectral properties.
1) Chromophore:- A chromophore is a part of a molecule responsible for absorbing UV or
visible light.
It typically contains π-electrons or non-bonding electrons (n) that can undergo electronic
transitions such as π → π*, or n → π*.
Absorption of light by chromophores promotes electrons from the ground state to an excited
state, resulting in an absorption band in the UV-visible spectrum.
● Examples of Chromophores:
Chromophore Common Structure Transition Type
C=C (Alkene) –CH=CH– π → π*
C=O (Carbonyl) –C=O n π*, π → π*
NO₂ (Nitro group) –NO₂ n π*, π → π*
Aromatic Rings Benzene-like structures π → π*
2) Auxochrome:- An auxochrome is a group that, when attached to a chromophore, does
not absorb light by itself in the visible region, but modifies the ability of the chromophore
to absorb light.
It usually contains lone pair electrons (e.g, –OH, –NH₂, –OR, –Cl).
● Auxochromes affect absorption in two main ways:
a) Shifting λmax (wavelength of maximum absorption)
b) Increasing intensity of absorption (ε value)
● Types of Shifts:
Shift Type Description. Caused by
Bathochromic shift. Absorption moves Electron-donating auxochromes
to longer
wavelength (red)
Hypsochromic shift. Absorption moves Electron-withdrawing auxochromes
to shorter
wavelength (blue)
