Principles of Ultraviolet-Visible Spectroscopy, Chemistry tutorial


The interaction between radiations and matter is an allure one. Most of the drug molecules absorb radiation in the ultraviolet region of the spectrum; however some are colored and therefore absorb radiation in the visible region. Absorption in the ultraviolet and visible regions of the electromagnetic spectrum corresponds to transitions between the electronic energy levels and gives helpful analytical information for both the organic and inorganic samples.

Absorption Spectroscopy:

The absorption of UV/Visible radiation takes place via excitation of electrons in the molecular structure to a higher energy state. Radiation is the form of energy and we are constantly reminded of its presence through our sense of sight and capability to feel the radiant heat. Radiation can be considered either as a continuous wave travelling via space, or as discrete photons of the similar energy. The wave approach is more helpful for lots of spectrometric approaches. It might be considered in terms of a wave motion where the wavelength, λ, is the distance between the two successive peaks (Figure shown below). The frequency, v, is the number of peaks passing a specific point per second. These terms are associated in the equation below:

c = vλ

Here, c is the velocity of light in a vacuum.

1019_Wavelength and Amplitude of a Wave.jpg

Fig: Wavelength and amplitude of a wave

The standard unit of wavelength is deduced in nanometers. Other units such as Angstrom and the mill micron (mµ) might be encountered; however their uses are now being discouraged.  In some situations, it is more convenient to consider light as a stream of particles termed as photons.

Photons are characterized by their energy 'E'. The energy of a photon is associated to the frequency of light via the equation below:

E = hv

Here, 'E' is the energy in joules (J), 'h' is Planck's constant, 6.626 x 10-34 Js, and 'v' is the frequency in inverse seconds (Hz).  From equations the above two equations, we can express that: 

E = hc/λ

From the above equations, we can observe that the energy of electromagnetic radiation is directly proportional to its frequency and inversely proportional to its wavelength. The electromagnetic radiation ranges from very low energy (that is, long wavelength, low frequency) radiation, such as radio waves and microwaves, to very high energy (that is, short wave-spectrum of interest to us as analytical chemists are illustrated in figure below. It is clear from the figure that the electromagnetic spectrum, to which the human eye responds, is just a very small part of all the radiant energy.

1771_Electromagnetic Spectrum range.jpg

Fig: Electromagnetic Spectrum range

Spectroscopy is basically the study of the interaction of light with matter. Matter is stated as materials comprised of molecules or atoms or ions. Whenever light strikes a sample of matter, the light might be absorbed by the sample, transmitted via the sample, reflected off the surface of the sample, or scattered through the sample. Samples can as well emit light after absorbing incident light; such a procedure is termed as luminescence.  

The ultraviolet (or UV) and visible region of the electromagnetic radiation covers the wavelength range from around 100 nm to around 800 nm. The vacuum ultraviolet region, which consists of the shortest wavelengths and highest energies (100-200 nm) is hard to make measurements in and is relatively uninformative. Helpful ultraviolet and visible absorption spectra are generated via the absorption of electromagnetic radiation with wavelengths in the 200 - 400 nm (UV) and 400 - 800 nm (Visible) regions of the electromagnetic radiation.

The atom comprises of a nucleus surrounded through electrons. Each and every element consists of a unique number of electrons, equivalent to its atomic number for a neutral atom of that element. The electrons are positioned in atomic orbital of different kinds and energies and the electronic energy states of atoms are quantized. The lowest energy, most stable electron configuration of an element is its ground state. The ground state is the normal electron configuration predicted from the 'rules' for filling a many-electron atom. Such rules are based on the position of the atom in the periodic table, the Aufbau principle, the Pauli Exclusion Principle and Hund's rule. For illustration, the ground state electronic configuration for sodium, atomic number 11, is 1s22s22p63s1 based on its position in the third row, first group of the periodic table and the requirement to account for 11 electrons. If energy of the right magnitude is given to an atom, the energy might be absorbed and an outer (that is, valence) electron promoted from the ground state orbital it is in, to a higher energy orbital. The atom is now in a higher energy, less stable, excited state. The electron will return spontaneously to the ground state, as the excited state is less stable than the ground state. In the process, the atom will emit energy; the energy will be equal in magnitude to the difference in energy levels between the ground and excited states.

The energy states related with molecules, like those of atoms are as well quantized. Whenever atoms join to form molecules, the individual atomic orbitals join to form a new set of molecular orbitals. Molecular orbital having electron density in the plane of the bonded nuclei, that is, all along the axis joining the bonded nuclei, are termed as sigma (σ) orbital. The molecular orbitals having electron density above and below the plane of the bonded nuclei are termed as pi (π) orbitals. Sigma and pi orbitals might be of two kinds, bonding or antibonding orbitals. As an illustration, the atomic orbitals of carbon, hydrogen and oxygen join in the molecule of propanone, C3H6O (figure shown below), therefore the three carbon atoms are joined in a chain by single (σ) bonds, the two outer carbons are each joined by σ bonds to three hydrogen atoms, whereas the central carbon atom is joined by a double bond to the oxygen atom, that is via both a σ bond and π bond. Moreover, the oxygen still consists of unpaired or non-bonded n electrons. This yields in a set of bonding and corresponding antibonding electronic orbitals or energy levels. Bonding energies are lower in energy as compare to the corresponding antibonding orbitals. Transitions might take place selectively between these levels, for illustration between π and π* levels.

Under normal conditions of pressure and temperature, the electrons in the molecule are in the ground state configuration, filling the lowest energy molecular orbitals available. The absorption of suitable radiant energy might cause an outer electron to be promoted to a higher energy excited state. As was the case by atoms, the radiant energy needed to cause electronic transitions in molecules lies in the visible and UV regions. The excited state of a molecule is less stable as compare to the ground state as with atoms; the molecules will spontaneously revert (or relax) to the ground state emitting UV or visible radiant energy. Dissimilar atoms, the energy states in molecules encompass rotational and vibrational sublevels; therefore whenever a molecule is excited electronically, there is frequently a simultaneous change in the vibrational and rotational energies. The net energy change is the sum of the electronic rotational and vibrational energy changes. Though, in the condensed states of solid and liquid, rotation is controlled.

The Organic molecules have carbon-carbon bonds, and bonds between carbon and other elements like hydrogen, oxygen, nitrogen, sulphur, phosphorus and halogens. Single bonds correspond to the bonding σ orbital, which consists of an associated antibonding σ* orbital.  Multiple bonds might as well be made up and correspond to the π bonding and π* antibonding orbitals. Bonding orbitals encompass lower energy, whereas antibonding orbitals encompass higher energy. Lone pair of electrons on atoms like oxygen is little changed in energy. Therefore, a molecule like propanone (acetone) has the structure shown below (figure illustrated below).

1013_Structure of Propanone.jpg

Fig: Structure of Propanone

Figure illustrates that the σ-σ* transitions need the largest energy change and take place at the lowest wavelengths, generally less than 190 nm, that is beneath the wavelengths measureable having most laboratory instrumentation. The π-π* transitions are very significant as they take place in all the molecules having multiple bonds and with conjugated structures, like aromatic compounds. The transitions take place around 200 nm, however the greater the extent of the conjugation, the closer the energy levels and the higher the noticed absorption wavelength.  Transitions comprising the lone pairs on heteroatom like oxygen or nitrogen might be n-σ*, which take place around 200 nm, or n-π*, which take place close to 300 nm. Such values are considerably modified by the specific structure and the presence of substituent (that is, auxochromes) in the molecules.

The single C-H and C-C bond associate to σ orbitals, the carbonyl double bond to the π orbitals and the unpaired electrons on the oxygen to the non-bonding n-levels. The energy levels might be grouped approximately as illustrated in the figure below. Transitions between σ and σ* levels and between π and π* are favored and those of the n electrons to the higher levels as well take place.

2188_Typical Transitions for Organic Molecules.jpg

Fig: Typical transitions for organic molecules

The energy levels comprised in transitions in the UV/visible region are the electronic levels of atoms and molecules. For illustration, however light atoms have widely space energy levels, a few heavy atoms have their outer orbitals close adequate altogether to provide transitions in the visible region. This accounts for the colours of iodides. Transition metals, having partially occupied d or f orbitals, often exhibit absorption bands in the visible region and these are influenced by the bonding of ligands example: Iron (III) reacts by thiocyanate ion to form an intense red colour due to the iron (III) thiocyanate complex, which might be employed to find out iron (III) in the presence of iron (II).

Instrumentation of Spectroscopy:

The components comprise (figure shown below):

1) The light sources: A deuterium lamp for the UV region from 190 - 350 nm and a quartz halogen or tungsten lamp for the visible region ranges from 350 - 900 nm.

2) The Monochromator: Employed to disperse the light into its component wavelengths that are further chosen by the slit.  The Monochromator is rotated in such a way that a range of wavelengths is passed via the sample as the instrument scans across the spectrum.

3) The optics: It might be designed to split the light beam in such a way that the beam passes via two sample compartments, and in such a double-beam instrument, a blank solution can then be employed in one compartment to correct the reading or spectrum of the sample. The blank is most generally the solvent in which the sample is dissolved.

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Fig: Schematic diagram of a UV/Visible Spectrophotometer

The UV spectrum is generally computer produced in modern UV spectrophotometers and it is characterized through wavy lines have a peak absorbance. Absorbance is plotted on the y-axis and wavelength on the x-axis (figure shown).

206_Representative UV Spectrum.jpg

Fig: Representative UV Spectrum (UV Spectrum of Ketoprofen)

Factors Governing Absorption of Radiation in the UV/Visible Region:

pH Effects:

pH will influence the structures of compounds having acidic and basic groups, and might cause considerable wavelength shifts.

Solvent Effects:

For π-π* transitions, the excited state is more polar as compare to the ground state, therefore it will tend to make dipole-dipole bonds having a polar solvent, like water or ethanol. This will lower the transition energy and increase the absorption peak wavelength. This is termed as the red shift (or a bathochromic shift). The tables of solvent corrections are available in the specialist texts. For n-π* transitions, the ground state is frequently more polar and might form hydrogen or dipole bonds by polar solvents. This raises the transition energy and lowers the peak wavelength shifts causing a blue shift (or hypsochromic shift).

Substituent Effects:

Substituents which modify the wavelength or absorptivity of a chromophore significantly are termed as auxochromes. The tables of the effect of substituent's plus rules for their application in specific structures are to be found in specialist texts. For illustration, an unsubstituted unsaturated ketone would encompass a peak maximum at around 215 nm. Substitution of a hydroxyl group on the carbon next to the carbonyl (α) increases the peak to 250 nm, and two alkyl groups on the next (β) carbons would lift it to 274 nm.

The table illustrated below lists some of the substituent effects of aromatic compounds. It must be noted that the phenoxide ion (-O-), that is present in alkaline solutions of phenols, absorbs at a considerably longer wavelength as compare to the parent phenol (-OH-). Usually electron donating and lone-pair substituents cause a red shift and more intense absorption. More complex shifts occur whenever there are more than one substituents present, and tables are given in the standard spectrometry texts listings these.  

Table: Absorption Maxima for Some Monosubstituted Benzenes Ph-R (in Methanol or Water) 

2298_Absorption Maxima for Monosubstituted Benzenes.jpg

Structure Effects:

The structure of organic molecules might be categorized in terms of the functional groups that they contain. Where such absorb UV or visible radiation in a specific region they are termed as chromophores. Some of the chromophores significant for analytical purposes are listed in the table shown below. This exhibits that the absorption via compounds having only σ bonds like hexane, or with lone pairs like ethanol will only take place beneath 200 nm. Such compounds are thus helpful solvents.

If more double bonds are present in the structure in conjugation (that is, two or more double bonds in a series separated through single bond), absorption occurs at longer wavelengths and having greater intensity. Such extended systems of double bonds are termed as 'chromophores'.

Table: UV Absorption characteristics of some Chromophores based on the Benzene Ring

276_UV Absorption Characteristics of Chromophores.jpg

A 1%, 1cm value provides a measure of the intensity of absorption. The most general chromophore found in drug molecules is a benzene ring. If the symmetry of the benzene ring is lowered via substitution, the bands in the benzene spectrum experience a bathochromic shift - a shift to longer wavelength. Substitution can comprise either extension of the chromophore or attachment of an auxochrome (that is, a group having one or more lone pair of electrons) to the ring and both. The hydroxyl group and amino group auxochromes are influenced through pH, undergoing bathochromic (moving to longer wavelength) and hyperchromic (that is, absorbing more strongly) shifts whenever a proton is removed under alkaline conditions, discharging an additional lone pair of electrons. The effect is most marked for aromatic amine groups.

Recognition of unknown organic samples can be considerably aided through considering the UV-visible absorption spectra. The given general rules might be employed as a guide.     


Possible Conclusion

a) No UV absorption presents Isolated double bonds.

Only σ bonds or lone pairs

b) Strong absorption between 200 and 250nm (ε ~ 1000) 

Aromatic rings

c) Weak absorption near 300 nm (ε ~ 1) 

Carbonyl compound

For illustration, an organic compound, C7H14O gave a UV spectrum having a peak at 296 nm and ε = 3.7 m2mol-1. Is it more probable to be a ketone or an alkene? The formula lets the possibility of merely one double bond. It should thus be an alkene having an isolated double bond, absorbing beneath 200 nm, or a ketone having a weak n-π* transition close to 300 nm. The value of both the absorption maximum and of the absorptivity recommends a ketone.

Beer-Lambert Law:

The measurement of light absorption through a solution of molecules is regulated by the Beer-Lambert Law. 

1758_Beer-Lambert Law.jpg

Fig: Beer-Lambert Law

Beer Lambert Law is represented as follows:

Log Io/It = A = εbc


Io = Intensity of the incident radiation 

It = Intensity of the transmitted radiation 

A = The absorbance and is a measure of the amount of light absorbed via the sample

ε = A constant recognized as the molar extinction coefficient and is the absorbance of a 1M solution of the analyte 

b = Path length of cell in cm, generally 1 cm

c = the concentration of analyte in moles litre-1.

Concentration and amounts are generally deduced in grams or milligrams instead of moles and therefore for the purposes of analysis of organic molecules, the Beer-Lambert equation is represented in the given form:

A = A (1%, 1cm) bc


A = Measured absorbance 

A (1%, 1cm) = Absorbance of a 1% w/v (1g/100ml) solution in a 1 cm cell 

b = Path length in cm (generally 1 cm)

c = Concentration of the sample in g/100 ml.  

As measurements are generally made in a 1 cm cell, the equation can be represented as:

C = [A/A (1%, 1cm)] which provides the concentration of the analyte in g/100 ml.

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