32 Ultraviolet-Visible (UV-Vis) Absorption Spectroscopy

Dr. Anchal Srivastava

Contents of this Unit

1.Introduction

2.  Basic Principle of UV-Vis Absorption Spectra 2.1 The Electromagntic spectrum

2.2 Radiation and the Atom

2.3 Radiation and the Molecule

2.4 Vibration and Rotation

3.  Principles of Absorption Spectroscopy: Beer’s and Lambert’s Law

4.  Construction and Working Principle

5.  Summary

Learning Outcomes

• After studying this module, you shall be able to
• Explain basic principles of UV-Visible spectroscopy Explain relevant terms of UV-Visible spectroscopy Describe the deviations of Beer’s Law.
• Explain working principles, taking spectrum and outline of atomic absorption spectroscopy device,

1. INTRODUCTION

Ultraviolet and visible spectrometers have been in general use for the last 35 years and over this period have become the most important analytical instrument in the modern day laboratory. In many applications other techniques could be employed but none rival UV-Visible spectrometry for its simplicity, versatility, speed, accuracy and cost-effectiveness.

2.  BASIC PRINCIPLE OF UV-Vis ABSORPTION SPECTRA 2.1 The Electromagnetic Spectrum

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 of Electromagnetic Spectrum

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 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 (ν), 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 (ν), which gives the number of waves per centimeter. Thus

1  cm = 10⁷ nm = 10⁸Ǻ

1  Ǻ = 10^-1 nm = 10^-8cm

The four quantities wavelength, wavenumber, frequency and velocity can be related to each other by following relationships

Wavelength (λ) = 1 / ν = c / ν

Wave-number (ν) = 1 / λ = ν / c

Frequency (ν) = c / λ = c ν

Velocity (c) = νλ = ν / ν

2.2 Radiation and The atom

Although it is convenient to describe electromagnetic radiation in terms of waves, it is necessary to define another model in order to demonstrate clearly the interactions that lead to selective absorption by an atom or molecule. A determining factor is the energy level of the radiation and it is therefore helpful to consider radiation as discrete packages of energy, or quanta. A quantum of light is known as a photon.

The absorption process depends upon an atomic structure in which each of the electrons of an atom has an energy level associated with its position in the atom. Permitted energy levels are finite and well defined, but an electron may be made to change to another level if a quantum of energy is delivered equal to the energy difference between the two levels. The original level is called the ground state and the induced level is known as the excited state. Excited states are generally unstable and the electron will rapidly revert to the ground state, losing the acquired energy in the process. Whilst the accepted model of atomic and molecular structure has arisen from the wave mechanical treatment of Schroedinger, it is convenient to employ an earlier model (that of Bohr) in order to explain more simply the electronic phenomena of interest in spectrophotometry.

The Bohr model defines an atom as having a number of electron shells, n1 – n2 – n3 etc, in which the increasing values of n represent higher energy levels and greater distance from the nucleus. Electrons rotate about the nucleus in orbits that may be characterised by the space they occupy and are designated s, p, d, etc according to their geometry. An atom may contain several electrons in multiple orbits in each shell (or each n level) but no orbit may contain more than two electrons (Fig. 3).

Figure.3. Diagrammatic presentation of simplified electron energy levels in an atom.

No two electrons can have identical energies but all can be assigned to groups corresponding to the shells, each of which has a clearly differentiated energy level. The effect of subjecting an atom to appropriate radiation is well demonstrated by considering atoms of sodium vapor.

A sodium atom at ground state (Na0) will absorb a photon at 589 nm to cause a transition of an electron in the outermost shell to a higher energy orbital.

Na₀ + 589 nm photon  Na_l

The same ground state atom will also absorb a 330 nm photon to promote a transition to its second excited state.

Na₀ + 330 nm photon  Na₂

Figure. 4 Absorption by gaseous atomic sodium

The diagram illustrates the higher energy required to cause a second level transition, corresponding with the higher energy (shorter wavelength) of the radiation that stimulates the transition (Fig. 4)

2.3 Radiation and the Molecule

Electronic Transitions

Electrons in the atom can be considered as occupying groups of roughly similar energy levels. In the more complicated molecular model, electrons associated with more than one nucleus, the so-called bonding electrons are particularly susceptible to energy level transitions under the stimulus of appropriate radiation.

Figure 5. Energy level diagram

The electrons concerned, usually p type electrons in the first or second shell, may be present in one of two conditions: (sigma) in localized bonds with a low probability of transitions (and therefore of absorption) or π (pi) where the transition probability is much higher. The presence of a carbon-carbon double bond in the molecule increases the likelihood of π type bonds, especially when conjugated double bonds are involved, i.e. double bonds that alternate with single bonds. The effect is still greater in the presence of nitrogen.

Chemical bonds are formed by overlapping atomic orbitals that result in molecular orbitals of one of three types: bonding (low energy), antibonding (high energy), or non-bonding. Energy absorption is most typically associated with transitions induced in electrons involved in bonding orbitals, and the atoms involved are, for the most part, those containing s + p electrons. Two types of bond must be mentioned:

(1)  bond with its related antibonding orbital designated  * and

(2)    π bonds with the corresponding π * antibonding orbital. The uninvolved n (non-bonding) electrons have

no antibonding orbital.

The full series of permitted electronic transitions (by UV/Vis absorption) is:

Figure 6. Schematic diagram of Electronic Transitions

An electron in a bonding s orbital is excited to the corresponding antibonding orbital. The energy required is large. For example, methane (which has only C-H bonds, and can only undergo *transitions) shows an absorbance maximum at 125 nm. Absorption maxima due to ® * transitions are not seen in typical UV-Vis. spectra (200 – 700 nm)

Saturated compounds containing atoms with lone pairs (non-bonding electrons) are capable of n * transitions. These transitions usually need less energy than * transitions. They can be initiated by light whose wavelength is in the range 150 – 250 nm. The number of organic functional groups with n * peaks in the UV region is small.

Most absorption spectroscopy of organic compounds is based on transitions of n or p electrons to the π * excited state. This is because the absorption peaks for these transitions fall in an experimentally convenient region of the spectrum (200 – 700 nm). These transitions need an unsaturated group in the molecule to provide the p electrons.

Molar  absorbtivities  from n π* transitions  are  relatively  low,  and  range  from  10  to100  L  mol-1 cm-1. π π* transitions normally give molar absorbtivities between 1000 and 10,000 L mol-1 cm-1 .

The solvent in which the absorbing species is dissolved also has an effect on the spectrum of the species. Peaks resulting from n π * transitions are shifted to shorter wavelengths (blue shift) with increasing solvent polarity. This arises from increased solvation of the lone pair, which lowers the energy of the n orbital. Often (but not always), the reverse (i.e. red shift) is seen for π π * transitions. This is caused by attractive polarization forces between the solvent and the absorber, which lower the energy levels of both the excited and unexcited states. This effect is greater for the excited state, and so the energy difference between the excited and unexcited states is slightly reduced – resulting in a small red shift. This effect also influences n π * transitions but is overshadowed by the blue shift resulting from solvation of lone pairs.

Charge – Transfer Absorption

Many inorganic species show charge-transfer absorption and are called charge-transfer complexes. For a complex to demonstrate charge-transfer behavior one of its components must have electron donating properties and another component must be able to accept electrons. Absorption of radiation then involves the transfer of an electron from the donor to an orbital associated with the acceptor.

2.4 Vibration and Rotation

The internal structure of a molecule may respond to radiant energy by more than just electronic transitions. In some molecules the bonding electrons also have natural resonant frequencies that give rise to molecular vibration while others exhibit a phenomenon known as rotation. Because the differences in energy levels associated with vibration and rotation are much smaller than those involved in electronic transitions, excitation will occur at correspondingly longer wavelengths.

Vibrational absorption is typically associated with the infrared region while the differences between energy levels related to molecular rotation are so small that far infrared or even microwave wavelengths are effective (Fig. 7)

Figure. 7 Relationship of wavelength and energy-induced transitions.

Because vibrational and rotational absorptions are primarily associated with spectral regions other than UV/Vis, it is necessary here to note only the effect on electronic absorption spectra. The principal effect is of “peak broadening”, i.e. the deviation of an observed absorption peak from the predicted shape.

For most absorbing species, especially in solution, absorption peaks do not appear as sharp lines at highly differentiated wavelengths, but rather as bands of absorption over a range of wavelengths. A principal reason is that an electronic transition is frequently accompanied by vibrational transitions between electronic levels (vibrational fine structure). In the same way each vibrational level may have associated rotational levels so that an absorption spectrum due to an electronic transition may well be a complex structure, with contributing components from vibrational and rotational absorption.

3.  PRINCIPLES OF ABSORPTION SPECTRSCOPY: BEER’s AND LAMBERT’s LAW

The greater the number of molecules that absorb light of a given wavelength, the greater the extent of light absorption and higher the peak intensity in absorption spectrum. If there are only a few molecules that absorb radiation, the total absorption of energy is less and consequently lower intensity peak is observed. This makes the basis of Beer-Lambert Law which states that the fraction of incident radiation absorbed is proportional to the number of absorbing molecules in its path.

When the radiation passes through a solution, the amount of light absorbed or transmitted is an exponential function of the molecular concentration of the solute and also a function of length of the path of radiation through the sample. Therefore,

-Log I / I = ε c l_o

Where I = Intensity of the incident light (or the light intensity passing through a reference cell)

o

I = Intensity of light transmitted through the sample solution

-1

c = concentration of the solute in mol l

ε   = molar absorptivity or the molar extinction coefficient of the substance whose light absorption is under investigation. It is a constant and is a characteristic of a given absorbing species (molecule or ion) in a particular solvent at a particular wavelength. ε is numerically equal to the absorbance of a solution of unit molar concentration (c = 1) in a cell of unit length ( l = 1) and its units are liters-moles-1 -cm-1. However, it is customary practice among organic chemists to omit the units.

The ratio I / I is known as transmittance T and the logarithm of the inverse ratio I / I is known as the absorbance

o                                                                                                                                    o

A. Therefore

– Log I_o / I = – log T = ε c l

and Log I_o / I = A = ε c l

or A = ε c l

For presenting the absorption characteristics of a spectrum, the positions of peaks are reported as λ (in nm)

max values and the absorptivity is expressed in parenthesis.

4.  CONSTRUCTION NAD WORKING PRINCIPLE

Figure 8. Schematic diagram of a double-beam UV-Vis. Spectrophotometer

Instruments for measuring the absorption of U.V. or visible radiation are made up of the following components;

1.      Sources (UV and visible)

2.      Wavelength selector (monochromator)

3.      Sample containers

4.      Detector

5.      Signal processor and readout

Instrumental components

Sources of UV radiation

It is important that the power of the radiation source does not change abruptly over it’s wavelength range.

The electrical excitation of deuterium or hydrogen at low pressure produces a continuous UV spectrum. The mechanism for this involves formation of an excited molecular species, which breaks up to give two atomic species and an ultraviolet photon. This can be shown as;

Both deuterium and hydrogen lamps emit radiation in the range 160 – 375 nm. Quartz windows must be used in these lamps, and quartz cuvettes must be used, because glass absorbs radiation of wavelengths less than 350 nm.

Sources of visible radiation

The tungsten filament lamp is commonly employed as a source of visible light. This type of lamp is used in the wavelength range of 350 – 2500 nm. The energy emitted by a tungsten filament lamp is proportional to the fourth power of the operating voltage. This means that for the energy output to be stable, the voltage to the lamp must be very stable indeed. Electronic voltage regulators or constant-voltage transformers are used to ensure this stability.

Tungsten/halogen lamps contain a small amount of iodine in a quartz “envelope” which also contains the tungsten filament. The iodine reacts with gaseous tungsten, formed by sublimation, producing the volatile compound WI2. When molecules of WI2 hit the filament they decompose, redepositing tungsten back on the filament. The lifetime of a tungsten/halogen lamp is approximately double that of an ordinary tungsten filament lamp. Tungsten/halogen lamps are very efficient, and their output extends well into the ultra-violet. They are used in many modern spectrophotometers.

Wavelength selector (monochromator)

• All monochromators contain the following component parts;
• An entrance slit
• A collimating lens
• A dispersing device (usually a prism or a grating) A focusing lens
• An exit slit

Polychromatic radiation (radiation of more than one wavelength) enters the monochromator through the entrance slit. The beam is collimated, and then strikes the dispersing element at an angle. The beam is split into its component wavelengths by the grating or prism. By moving the dispersing element or the exit slit, radiation of only a particular wavelength leaves the monochromator through the exit slit.

Figure 9: Czerney-Turner grating monochromator

Cuvettes

The containers for the sample and reference solution must be transparent to the radiation which will pass through them. Quartz or fused silica cuvettes are required for spectroscopy in the UV region. These cells are also transparent in the visible region. Silicate glasses can be used for the manufacture of cuvettes for use between 350 and 2000 nm.

Detectors

The photomultiplier tube is a commonly used detector in UV-Vis spectroscopy. It consists of a photoemissive cathode (a cathode which emits electrons when struck by photons of radiation), several dynodes (which emit several electrons for each electron striking them) and an anode.

A photon of radiation entering the tube strikes the cathode, causing the emission of several electrons. These electrons are accelerated towards the first dynode (which is 90V more positive than the cathode). The electrons strike the first dynode, causing the emission of several electrons for each incident electron. These electrons are then accelerated towards the second dynode, to produce more electrons which are accelerated towards dynode three and so on. Eventually, the electrons are collected at the anode. By this time, each original photon has produced 106 – 107 electrons. The resulting current is amplified and measured.

Photomultipliers are very sensitive to UV and visible radiation. They have fast response times. Intense light damages photomultipliers; they are limited to measuring low power radiation.

Figure 9: Cross section of a photomultiplier tube

The linear photodiode array is an example of a multichannel photon detector. These detectors are capable of measuring all elements of a beam of dispersed radiation simultaneously.

A linear photodiode array comprises many small silicon photodiodes formed on a single silicon chip. There can be between 64 to 4096 sensor elements on a chip, the most common being 1024 photodiodes. For each diode, there is also a storage capacitor and a switch. The individual diode-capacitor circuits can be sequentially scanned.

In use, the photodiode array is positioned at the focal plane of the monochromator (after the dispersing element) such that the spectrum falls on the diode array. They are useful for recording UV-Vis. absorption spectra of samples that are rapidly passing through a sample flow cell, such as in an HPLC detector.

Charge-Coupled Devices (CCDs) are similar to diode array detectors, but instead of diodes, they consist of an array of photocapacitors.

5. SUMMARY

Ultraviolet-visible (UV-vis) radiation has the property of being able to promote an electron to a higher energy level.

The Beer-Lambert law is used to relate absorption of UV-vis radiation to the concentration of the substance, the path length of the cell and the molar absorption coefficient. The molar absorption coefficient is a constant for every substance and is a measure of the amount of radiation absorbed per unit concentration of a substance

Molecules containing π-electrons or non-bonding electrons (n-electrons) can absorb the energy in the form of ultraviolet or visible light to excite these electrons to higher anti-bonding molecular orbitals. The more easily excited the electrons, the longer the wavelength of light it can absorb. Based on the fact of four type of transition- π-π*, n-π*, σ-σ*, and n-σ*. The energy required for various transitions obey the following order σ-σ* > n-σ* > π-π* > n-π*.