Spectroscopy may be defined as the study of interaction between electromagnetic radiations and matter.

Transcription

1 Spectroscopy may be defined as the study of interaction between electromagnetic radiations and matter. Spectroscopy has a wide range of applications. It is heavily used in astronomy and remote sensing. Most large telescopes have spectrographs which are used either to measure the chemical composition and physical properties of astronomical objects or to measure their velocities (from the doppler shift of the spectral lines). In medical diagnosis, various techniques employed (such as MRI, EEG, X-ray etc.) are based on spectroscopy. It has a wide range of applications in the study of atomic and molecular structure. It is an important tool in the hands of analytical chemists for the identification and quantitative analysis of chemical substances. As we shall be concerned with what spectroscopy can tell us of the structure of matter, so it is essential to discuss briefly the nature of electromagnetic radiation and the sort of interactions which may occur. Electromagnetic Radiation Electromagnetic radiation is a transverse energy wave that is composed of an oscillating electric field component, E, and an oscillating magnetic field component, M. The electric and magnetic fields are orthogonal to each other, and they are orthogonal to the direction of propogation of the wave. A wave is described by the wavelength,, which is the physical length of one complete oscillation, and the frequency,, which is the number of oscillations per second. The figure shows one wavelength of a wave of light. Schematic of an electromagnetic wave Electromagnetic radiation, of which visible light forms an obvious but very small part, may be considered as a simple harmonic wave, propagating from a source and traveling in a straight line (except when refracted or reflected). This can be represented by the sine wave equation y = A sin 2πx/λ where λ is wavelength of the wave. Energy associated with wave is given by the relation E = hν where ν is the frequency (units: s -1 or hertz) of the wave and is given by ν = c/λ; h is Planck s constant (h = 6.63 x j s). Since λ = 1/ ν; ν = c ν where ν is the wave number (units: cm -1 ). The names we give electromagnetic radiation for different wavelength and frequency ranges are listed in the order of their energies and collectively called electromagnetic spectrum. Electromagnetic Spectrum Electromagnetic spectrum constitutes all possible electromagnetic radiations arranged according to the order of their increasing (or decreasing) energies. As already mentioned 1

2 electromagnetic radiation at a particular wavelength λ has an associated frequency ν and energy E. Thus electromagnetic spectrum may be expressed equally well in terms of all the three quantities. In the increasing order of frequency (and hence energy) the regions of electromagnetic spectrum are as follows. 1. Radiofrequency region: Electromagnetic waves corresponding to frequencies 3 x x Hz or wave length 10 cm to 1 cm lie in this region. This amount of energy change arises from change in spin of electron or nucleus and is of the order of to 10 J/mol. 2. Microwave region: Electromagnetic waves corresponding to frequencies 3 x x Hz or wavelength 1 cm to 100 µm lie in this region. Separations between rotational energy levels of molecules are of the order of hundreds of joules per mol and a transition of an electron from one rotational level to the other is accompanied by absorption of energy in this region. 3. Infrared region: Electromagnetic waves corresponding to frequencies 3 x x Hz or 100 µm to 1 µm lie in this region. These radiations (of corresponding energy of the order 10 4 J/mol) are absorbed when transition takes place between vibrational levels. The study of such absorptions is called vibrational spectroscopy. 4. Visible UV regions. Radiations corresponding to energy changes of the order hundreds of kilo joules (ν = 3 x x Hz or λ= 1 µm to 10 nm). Absorptions in this region occur due to transition of valence electrons between electronic energy levels. So, study of these transitions is termed as electronic spectroscopy. Visible region lies between 400 nm to 800 nm whereas the rest is UV region. 5. X ray region: Energy of these radiations fall between ten thousands of kj(ν = 3 x x Hz or λ= 10 nm to 100 pm). These radiations are emitted or absorbed when inner electrons of an atom or a molecule change energy. 6. γ- ray region: γ radiations are associated with energy equal to J/g atom (ν = 3 x x Hz or λ= 10 pm or less). Such a high energy change occurs due to rearrangement of nuclear particles. Cosmic rays possess energy even higher to γ rays but are found only in outer space. Type of Radiation gamma-rays X-rays ultraviolet visible near-infrared infrared microwaves Type of Transition nuclear inner electron outer electron outer electron outer electron molecular vibrations molecular vibrations molecular rotations, electron spin flips* 2

3 radio waves nuclear spin flips* Given below is a representation of entire electromagnetic spectrum. The visible light constitutes only a small portion of entire spectrum and is shown separately. The energy E of an electromagnetic radiation is quantized. In 1900, Max Planck gave the idea of quantization of energy. That is, the energy of an oscillator is discontinuous and any change in its energy can occur only by means of a jump between two distinct energy states. The idea was later extended to many other forms of energy of matter. This implies that a transition can take place between two distinct energy levels associated with an atom or molecule and energy equal to the difference between these two levels E can only be absorbed or emitted. If we take a molecule or atom in state 1 associated with energy E 1 and direct on to it a beam of radiation of a single frequency ν (monochromatic radiation), the energy will be absorbed from the beam of radiation and atom or molecule will jump from state 1 to state 2. A detector placed to collect the radiation after its interaction with the species will show that intensity of the radiation has decreased. If we use a beam containing a wide range of frequencies, the detector will show that energy has been absorbed only from the frequency ν = E/h, all other frequencies remain undiminished in intensity. E is the difference in the energies of level 1 and level 2. In this way we have produced an absorption spectrum. Alternatively, the species may already be in energy state 2 and may revert back to the state 1 with emission of radiation. The detector would show this radiation to have frequency ν = E/h. The emission spectrum so obtained is complementary to the absorption spectrum discussed above. Spectrum: The data obtained from spectroscopy is called a spectrum. A spectrum is a plot of intensity of energy detected versus wavelength (or frequency or wave number or mass or momentum etc.) 3

4 A spectrum can be used to obtain information about atomic and molecular energy levels, molecular geometries, chemical bonds, interactions of molecules, and related processes. ften, spectra are used to identify the components of a sample (qualitative analysis). Spectra may also be used to measure the amount of material in a sample (quantitative analysis). Atomic spectroscopy deals with the interaction of electromagnetic radiations with atoms. As explained above, both absorption and emission atomic spectra can be observed for atoms. Thus, atomic spectroscopy can be categorized into two types atomic absorption spectroscopy and atomic emission spectroscopy. In atomic absorption spectroscopy, the atoms are in their ground states initially and can be excited by irradiating them with radiations of appropriate frequency as given by the energy difference of their ground and excited energy levels. These energy levels are highest occupied valence orbitals and lowest unoccupied orbitals respectively. The excitation of valence electrons takes place during the transition. During atomic emission spectroscopy the electrons are already in excited state and revert back to the ground state with emission of radiation of frequency equivalent to the energy difference of their ground and excited energy levels. An important application of atomic spectroscopy is Flame absorption and Flame emission spectroscopy. The flame is used to bring the sample solution to gaseous state in former and excite the atoms to higher state besides bringing the sample solution to gaseous state for later. These techniques are used for detection of trace elements. For example Flame absorption spectra are used to detect Mg in water, V in lubricating oils, Cd, Cu, Zn, Ni, As etc. in soil, Sn in canned fruit juices and flame emission spectra are obtained for analysis of alkali metals in biological fluids or tissues. Molecular spectroscopy deals with the interaction of electromagnetic radiations with molecules. As compared to atomic spectroscopy, it is relatively complex. Before discussing different types of molecular spectroscopy, let us discuss the kind of energies a molecule possesses. The energy possessed by a polyatomic molecule may be due to contribution from translational energy (arising from change of its centre of gravity as a result of motion), viberational energy (arising from to and fro motion of the nuclei of the molecule so that its centre of gravity does not change) and rotational energy (arising from its rotation about an axis perpendicular to the internuclear axis and passing through the centre of gravity of the molecule). Another form of energy that a molecule possesses is the electronic energy, associated with the transition of an electron from the ground state to excited state by absorbing photon of suitable energy. As an approximation, suggested by Max Born and J. Robert ppenhiemer, nuclear motions can be separated from electronic one because the nuclei being massive move much slowly compared to electrons and thus may be regarded as stationary. Hence various molecular energies can be estimated by solving Shrodinger wave equation by fixing position of nuclei. Therefore change in translational energy can be ignored while calculating the total molecular energy. So, the total energy of a molecule is considered as 4

5 the sum of electronic, viberational and rotational energies associated with corresponding electronic, viberational and rotational levels. E = E electronic + E viberational + E rotational Associated with each electronic level is a number of viberational level and with each viberational level a series of rotational levels are associated. The order of energies of these levels are E electronic > E viberational > E rotational. All these energies are quantized. The order of magnitude of these energy changes are E electronic E viberational X 10 3 E rotational X Electronic Spectroscopy Electronic spectroscopy deals with the excitation of valence electrons of the molecules or species from the lower electronic level to higher electronic level. The energy thus absorbed lies in the ultraviolet or visible region of the electromagnetic spectrum. The UV region extends from nm whereas the visible region extends from nm. Below 200 nm, 2 absorbs strongly. Therefore for all practical purposes, wavelength range of nm is chosen. Electronic transitions When an atom or molecule absorbs energy, electrons are promoted from their ground state to an excited state i.e. from the outermost filled orbital (HM) to lowest in energy empty orbital (LUM). There are three types of electronic transitions which can be considered; 1. Transitions involving π, σ, and n electrons 2. Transitions involving charge-transfer electrons 3. Transitions involving d and f electrons Absorbing species containing π, σ, and n electrons Absorption of ultraviolet and visible radiation in organic molecules occurs due to excitation of electrons from bonding or non bonding orbitals (ground state) to the antibonding molecular orbitals (excited state) which are usually vacant when the molecule is in ground state. The electrons present in π, σ, and n orbitals are called π, σ, and n electrons. The antibonding orbital associated with σ electrons is σ, with π is π* while n (non bonding) electrons are not associated with any antibonding molecular orbital since they are not involved in bond formation. Consequently, n electrons are present in the atomic orbitals which have higher energy as compared to π electrons. 5

6 Possible electronic transitions are (i) σ σ * transitions (ii) n σ * transitions (iii) n π * transitions (iv) π π * transitions σ σ * Transitions An electron in a bonding σ orbital is excited to the corresponding antibonding orbital. The energy required is large. These transitions are shown only by those compounds in which all the electrons are involved in the σ bond formation, e. g. saturated hydrocarbons. For example, methane (which has only C-H bonds, and can only undergo σ σ * transitions) shows an absorbance maximum at 125 nm. Absorption maxima due to σ σ * transitions are not seen in typical UV-Visible spectra ( nm) n σ * Transitions Transitions in which non bonding electrons are excited to σ * antibonding molecular orbitals are called n σ * transitions. Saturated compounds containing atoms with lone pairs (non-bonding electrons) are capable of n σ * transitions. For example, alcohol, ethers, amines, alkyl halide etc. These transitions usually need less energy than σ σ * transitions. n π * Transitions When electrons present in non bonding orbital are excited to π * orbitals, the transitions are called n π * transitions. All organic compounds containing double bond between carbon and some heteroatom (, S, N etc.) show this transition. π π * Transitions Transitions in which a π electron is excited to π orbital are called π π * transitions. This type of transition occurs in unsaturated compounds (alkenes, alkynes, aromatic compounds). Most absorption spectroscopy of organic compounds is based on transitions of n or π electrons to the π * excited state. This is because the absorption peaks for these transitions 6

7 fall in an experimentally convenient region of the spectrum ( nm). These transitions need an unsaturated group in the molecule. Transitions involving charge-transfer electrons These transitions occur mainly in inorganic complexes formed by coordination of ligands to metal atom. The transitions involve electron transfer from one part of the complex to another. More specifically, an electron moves from an orbital that is mainly ligand in character to one that is mainly metal in character or vice versa. These are intense absorptions. Some examples of these absorptions are permanganate ion, dark colour of I 2 benzene solution. Transitions involving d and f electrons The electronic spectra of transition metals are referred to as d-d transitions because they involve orbitals which are mainly metal-d in character. The spacing between the energy levels depends upon the geometry of the complex, nature of the ligand, oxidation state of metal etc. Similarly, the absorption spectra of lanthanides and actinides result from f-f transitions. Peak Shape: UV and visible spectra of molecules consist of bands rather than peaks. In a molecule, the atoms can rotate and vibrate with respect to each other. These vibrations and rotations also have discrete energy levels, which can be considered as being packed on top of each electronic level as shown below. Therefore in molecules, each electronic state is associated with a number of closely spaced vibrational sub states and each vibrational state has a series of closely spaced rotational levels. In principle, electronic excitations can occur from any vibrational sublevel of the ground electronic state but these generally occur from the ν o level of the ground electronic state because this is the highly populated state for molecules at room temperature. As a result of these transitions, the absorption spectrum of the molecule in vapour state would contain a large number of closely spaced lines which can only be separated with the help of instruments of high resolution. However if the spectrum is taken in solution state, this fine structure is generally not observed because of solvent-solute interactions which tend to smooth out the spectrum into bands. Thus UV and visible spectra of molecules consist of bands rather than peaks. 7

8 Franck Condon principle As stated above the width of electronic absorption bands in liquid samples is due to their unresolved vibrational structure. This vibrational structure can be resolved in gases. The appearance of the vibrational pattern can be explained on the basis of Franck Condon principle. This principle states that because the nuclei are so much more massive than electrons, an electronic transition takes place faster than the nuclei can respond. So nuclear framework remains constant during this excitation and we imagine the transition as being represented by the vertical line. During an electronic transition, a change from one vibrational energy level to another will be more likely to happen if the two vibrational wave functions overlap more significantly. This can also be explained as follows. The absorption of a photon is a practically instantaneous process, since it involves only the rearrangement of practically inertia-free electrons. James Frank recognized the obvious: the nuclei are enormously heavy as compared to the electrons. Thus, during light absorption, that occurs in femtoseconds, electrons can move, not the nuclei. The much heavier atomic nuclei have no time to readjust themselves during the absorption act, but have to do it after it is over, and this readjustment brings them into vibrations. This is best illustrated by potential energy diagrams, such as that shown below. It is an expanded energy level diagram, with the abscissa acquiring the meaning of distance between the nuclei, rxy. The two potential curves show the potential energy of the molecule as a function of this distance for two electronic states, a ground state and an excited state. Excitation is represented, according to the Franck-Condon principle, by a vertical arrow (A). This arrow hits the upper curve, except for very special cases, not in its lowest point, corresponding to a non vibrating state, but somewhere higher. This means that the molecule finds itself, after the absorption act, in a nonequilibrium state and begins to vibrate like a spring. 8

9 Selection Rules for electronic spectroscopy Both atomic and molecular spectra result from the transition between different energy levels. These transitions take place between certain specific energy levels and do not take place between any two energy levels. Thus only restricted transitions can take place. The restrictions thus applied on transitions are called selection rules. Transition permitted by the selection rules is known as allowed transition where as that not permitted is called a forbidden transition. For a molecular electronic transition, the selection rules are as follows. 1. Transition which does not involve change in spin quantum number is allowed transition i.e. S = 0 i.e. singlet-triplet transitions are forbidden. 2. nly one electron is involved in a transition. 3. The transitions between orbitals of different symmetry can only occur (g u and u g allowed; but g g and u u not allowed). The allowed transitions have high value of ε max (extinction coefficient), generally more than An example of allowed transition is π π * transition. However forbidden transitions can also take place, but its probability is low. ε max for such transitions lies between The n π * transitions are forbidden but give rise to weak absorption. For example, benzophenon shows two absorption bands in UV spectrum. (i) λ max 252 nm, ε max 20,000 (allowed) (ii) λ max 325 nm, ε max 180 (forbidden) In π π * transition, both π and π * orbitals lie in the same plane and hence overlap between these orbitals in excited state is quite large. As a result π π * transitions are highly probable. n the other hand, n π * transitions are symmetry forbidden because n electrons on the heteroatom are in a plane perpendicular to the π * M. Since the regions of space of these two orbitals overlap poorly, the probability of excitation of n electron to π * orbital is low. Though symmetry forbidden, this transition occurs due to viberational interaction and twisting of molecule in excited state, it becomes partially allowed. Beer-Lambert law Many compounds absorb ultraviolet (UV) or visible (Vis.) light. The diagram below shows a beam of monochromatic radiation of intensity I 0, directed at a sample solution. Absorption takes place and the beam of radiation leaving the sample has intensity I t. Diagram of Beer-Lambert absorption of a beam of light as it travels through a cuvette of size l. 9

10 Lambert investigated the relationship between the intensity of incident light I 0 and that of transmitted light I t. Lambert law states that The rate at which the intensity of light decreases with thickness of the absorbing medium is proportional to the intensity of the incident light. -di/dx = ki or di/i = -kdx where I is the intensity of incident light, x is the thickness of absorbing medium and k is a constant called absorption coefficient of the medium for light of a particular wavelength. The negative sign indicates that I decreases as x increases. n integrating the above equation between I = I o (incident light) to I = I t (transmitted light), we get I t x=l di/i = - kdx I o x=o ln I t /I o = -kl Taking antilog I t /I o = e -kl or I t = I o e -kl Beer s law states that equal fractions of the incident radiations are absorbed by layers of solution with equal concentrations and same thickness. n combining the two laws, Beer-Lambert law states n passing through a solution, rate of decrease of intensity of incident monochromatic radiation with thickness of the solution is proportional to the intensity of incident light and concentration of the solution. So, -di/dx = k Ic where c is the molar concentration of the solution, k is molar absorption coefficient. n integrating, the equation between I = I o (incident light) to I = I t (transmitted light), we get ln I t /I o = -k lc or I t /I o = e -k lc I t = I o e -k l c I t = I o. 10 -k lc/2.303 Substituting k /2.303 = ε, we get I t = I o. 10 -εlc ε is now absorption coefficient or molar absorbtivity with units of L mol -1 cm -1. The equation can be written as - log I t /I o = εcl or log I o /I t = εcl Put log I o /I t = A So, A= εcl A is the absorbance or optical density and has no units. The molar absorptivity or molar extinction coefficient ε of a chemical species at a given wavelength is a measure of how strongly the species absorbs light at that wavelength. It is an intrinsic property of the species; the actual absorbance of a sample is dependent on its thickness l and the concentration c of the species. The amount of radiation absorbed may be measured as Transmittance. Transmittance is defined as the intensity of transmitted light to that of incident light. 10

11 Transmittance, T = I t / I 0 % Transmittance, %T = 100 T The relationship between absorbance and transmittance is illustrated in the following diagram. So, if all the light passes through a solution without any absorption, then absorbance is zero, and percent transmittance is 100%. If all the light is absorbed, then percent transmittance is zero, and absorption is infinite. A = εbc tells us that absorbance depends on the total quantity of the absorbing compound in the light path through the cuvette. If we plot absorbance against concentration, we get a straight line passing through the origin (0,0). Note that the Law is not obeyed at high concentrations. This deviation from the Law is not dealt with here. Numerical Problems on Beer-Lambert law 1. What is molar extinction coefficient of a solute which absorbs 90% of a certain wavelength of light beam passed through 1 cm cell containing 0.25 M solution? 2. Light of definite wavelength was passed through a cell of 4 cm thickness containing 0.02 M solution of a given substance. If molar extinction coefficient is 10.0 litre -1 cm -1, calculate the optical density and percentage transmittance. 3. Determine the transmittance, absorbance and molar extinction coefficient of a solution which transmits 50% of a monochromatic light when passed through 1 cm thick cell containing 5 X 10-3 M of solute. Ultraviolet-visible spectrophotometer The instrument used in Ultraviolet-visible spectroscopy is called ultraviolet-visible spectrophotometer. To obtain absorption information, a sample is placed in the spectrophotometer and ultraviolet or visible light at a certain wavelength, or range of 11

12 wavelengths, is transmitted through the sample. The spectrophotometer measures how much of the light is absorbed by the sample. The intensity of light before going into a certain sample is symbolized by I 0. The intensity of light remaining after it has gone through the sample is symbolized by I. The fraction of light transmittance is (I / I 0 ), which is usually expressed as a percent Transmittance (%T). From this information, the absorbance of the sample is determined for that wavelength or as a function for a range of wavelengths. Sophisticated UV/ Vis spectrophotometers often do this automatically. Although the samples could be solid, or even gaseous, they are usually liquid. A transparent cell, often called a cuvette, is used to hold a liquid sample in the spectrophotometer. The pathlength L through the sample is then the width of the cell through which the light passes through. Simple, economical spectrophotometers may use cuvettes shaped like cylindrical test tubes, but more sophisticated ones use rectangular cuvettes, commonly 1 cm in width. For just visible spectroscopy, ordinary glass cuvettes may be used, but ultraviolet spectroscopy requires special cuvettes made of an ultraviolet-transparent material such as quartz. A diagram of the components of a typical spectrometer is shown in the following diagram. The functioning of this instrument is relatively straightforward. A beam of light from a visible and/or UV light source is separated into its component wavelengths by a prism or diffraction grating. Each monochromatic (single wavelength) beam in turn is split into two equal intensity beams by a half-mirrored device. ne beam, the sample beam, passes through a small transparent container (cuvette) containing a solution of the compound being studied in a transparent solvent. The other beam, the reference, passes through an identical cuvette containing only the solvent. The intensities of these light beams are then measured by electronic detectors and compared. The intensity of the reference beam, which should have suffered little or no light absorption, is defined as I 0. The intensity of the sample beam is defined as I. ver a short period of time, the spectrometer automatically scans all the component wavelengths in the manner described. The ultraviolet (UV) region scanned is normally from 200 to 400 nm, and the visible portion is from 400 to 800nm. 12

13 Ultraviolet-visible spectrum An ultraviolet-visible spectrum is essentially a graph of light absorbance versus wavelength in a range of ultraviolet or visible regions as shown. If the sample compound does not absorb light of of a given wavelength, I = I 0. However, if the sample compound absorbs light then I is less than I 0, and this difference may be plotted on a graph versus wavelength, as shown on the right. Absorption may be presented as transmittance (T = I/I 0 ) or absorbance (A= log I 0 /I). If no absorption has occurred, T = 1.0 and A= 0. Most spectrometers display absorbance on the vertical axis, and the commonly observed range is from 0 (100% transmittance) to 2 (1% transmittance). The wavelength of maximum absorbance is a characteristic value, designated as λ max. Different compounds may have very different absorption maxima and absorbances. Intensely absorbing compounds must be examined in dilute solution, so that significant light energy is received by the detector, and this requires the use of completely transparent (non-absorbing) solvents. The most commonly used solvents are water, ethanol, hexane and cyclohexane. Solvents having double or triple bonds, or heavy atoms (e.g. S, Br & I) are generally avoided. Because the absorbance of a sample will be proportional to its molar concentration in the sample cuvette, a corrected absorption value known as the molar absorptivity is used when comparing the spectra of different compounds. Chromophores The term chromophore was originally applied to a system which imparts colour to a compound (Greek chromophorus = colour carrier). In UV- Visible spectroscopy, the term is used in a broader sense. A functional group that absorbs radiations in UV Visible region, irrespective of the fact that whether it imparts colour to the compound or not, is termed as chromophore. For example, a carbonyl group is a chromophore because it absorbs UV radiations. Certain chromophores and their λ max and ε max are listed below. Chromophore λ max ε max >C=C< 175 (π π * ) 15,000 -C C- 170 (π π * ) 10,000 >C= 165 (n σ * ), 190 (π π * ) 280 (n π * ) 5,000, 16,000 15, 13

14 -CH 208 (n π * ) 32 -CNH (n π * ) 63 -CR 211 (n π * ) 57 -N (n σ * ), 274 (n π * ) 5,000, 17 C N 165 (n π * ) 65 N=N 338 (n π * ) 4 Auxochromes Groups which themselves do not show any characteristic absorption above 200 nm but when attached to a given chromophore, cause a shift of the absorption towards longer wavelength along with increase in intensity of the absorption. For example: - H, -NH 2, -SH, -R, -NHR, -NR 2, -SR, halogens etc. For example, benzene shows an absorption band at λ max 254 nm with ε max of 230. But aniline shows λ max at 280 with ε max of This is due to conjugation between nitrogen nonbonding electrons and benzene π electrons. However, if a proton gets attached to aniline and forms NH 3 + then the absorption of resulting anilinium ion is observed at 254 nm because now nitrogen s lone pair is not conjugating the ring π electron density. Another important fact to remember is that auxochromes shift the absorption of π π * transitions towards longer wavelength whereas reverse effect is observed for n π * transitions (as discussed in following sections). NH 2 + NH3 λ max 254 nm 280 nm 254 nm ε max H 3 C H 3 C H H 3 C NH 2 H 3 C Cl λ max Fig. 1 14

15 The absorption band is affected in four different ways. 1. Bathochromic shift or Red shift 2. Hypsochromic shift or Blue shift 3. Hyperchromic effect 4. Hypochromic effect Bathochromic shift or Red shift The shift of absorption maxima to longer wavelength is called bathochromic shift or red shift. This can be achieved by (i) attaching an auxochrome to a carbon carbon double bond or to benzene ring as discussed above (ii) conjugation of two chromophores (iii) decreasing of polarity of the solvent. All these factors bring about decrease in energy difference between the orbitals involved in the transition. Hypsochromic shift or Blue shift A shift in absorption maxima to shorter wavelength is called Hypsochromic shift or Blue shift. This can be achieved by (i) Attaching auxochrome to the C= (carbonyl) group. For example, n π * transition of acetaldehyde appears at 293 nm (ε = 12) but it shifts to 235 (ε = 53) if H is replaced by Cl. Similarly, λ max for acetamide, ethyl acetate and acetic acid appear at 214 nm (ε = 63), 204 (ε = 60) and 204 (ε = 32) respectively (Fig. 1). This may be due to more decrease in the energy of n orbitals compared to π * orbitals, caused by I effect of heteroatoms (, N, Cl). These atoms withdraw electrons from C= thus causing lone pair of carbonyl to be held more strongly. (ii) Increasing the polarity of the solvent. For example, n π * transition of acetone in hexane appears at 279 nm and at 264 nm in water. This is because H bonding lowers the energy of n electrons and thus increases the energy gap between the two orbitals. Hyperchromic effect This effect leads to an increase in molar absorptivity. Red shift is generally accompanied by hyperchromic effect. Hypochromic effect This effect leads to decrease in molar absorptivity. Blue shift is generally accompanied by hypochromic effect. Solvent Effects The polarity of a solvent usually shifts the position of an absorption band towards longer or shorter wavelength. For different type of transition, the effect is different. (i) n π * and n σ * transitions generally shift towards lower wavelength if polar solvent is used. This is because the n electrons are relatively more stabilized in the ground state due to hydrogen bonding or dipole - dipole interactions with polar solvent than in excited state. As a result, the energy difference between ground and excited state increases resulting in shift towards shorter wavelength. It is to be noted that greater the shift, stronger the H bonding. The extent of H bonding decreases in the order: water > methanol > ethanol > chloroform. λ max for n π * transition of acetone appears at 279 nm in hexane (non polar) and shifts to shorter wavelength as the polarity of solvent increases. Solvent CHCl 3 C 2 H 5 H H H 2 λ max (nm)

16 (ii) π π * transitions shift to longer wavelength as the polarity of solvent increases. This is due to the fact that π * state, being more polar in nature than π state, (which is almost non polar because electrons are evenly distributed on two atoms involved in π bond formation) is more stabilized on interaction with polar solvent than the ground state. Effect of Conjugation When the two chromophoric groups are conjugated, a large effect on the spectrum is produced because the π electron system is now spread over at least four atomic centers. As a result, the high intensity absorption band due to π π transition is shifted by nm towards longer wavelength with increase in intensity (red shift and hyperchromic effect). Similarly, if carbonyl group conjugates with double bond, absorption bands due to both π π and n π transition shift to longer wavelengths. Hence conjugation produces red shift and hyperchromic effect. UV spectra of Dienes Alkenes such as ethylene absorb at 175 nm (ε = 15,000). But when two double bonds conjugate as in 1, 3 butadiene the absorption band shifts to 217 nm (ε = 21,000). This may be explained as follows. According to molecular orbital theory atomic p orbital on each of the carbon atoms combine to make two π molecular orbitals (Fig. 7.11). The two electrons forming π bond are filled in the lowest orbitals designated as ψ1 in the Fig. Now consider 1, 3- butadiene, which has four atomic p orbitals and four p electrons (each contributed by each one of the carbon atoms of butadiene) that form π system of two conjugated double bonds. The four atomic orbitals combine to form four molecular orbitals as shown in Fig The highest occupied M (HM) here is designated as ψ2 as shown in the fig. The energy difference from ψ2 (HM) to ψ3* (LUM) is lesser in energy compared to that of ψ1 and ψ2* of ethylene. The transition for diene is thus involves lesser energy compared to that of ethylene (red shift). More over greater the number of conjugated double bonds, greater is the red shift. Moreover the energy difference between HM and LUM becomes progressively smaller with increasing conjugation The effect is similarly observed in homoannular dienes (two double bonds in a ring) and heteroannular dienes (conjugating double bonds in different rings). 16

17 Woodward Fieser Rules for calculation of λ max of π π * transitions of simple conjugated dienes. Woodward formulated certain empirical rules for calculating λ max for π π * transitions of dienes and polyenes which were later modified by Fieser. According to these rules each diene has a certain fixed basic value, which is altered by the substituents in the following manner. Parent heteroannular or open chain diene 214 nm Parent homoannular diene 253 nm For each double bond extending conjugation add 30 nm 17

18 For each alkyl substituent or ring add 5 nm For each exocyclic double bond add 5 nm -C add 0 nm -R add 6 nm -SR add 30 nm -NR 2 add 60 nm -Cl, -Br add 5 nm UV Spectra of unsaturated ketones: enones The UV spectra of unsaturated carbonyl compounds in which the C=C and C= are not conjugated are simple summation of the absorption of the ethylene and carbonyl chromophores. In ethylene the absorption occurs at far UV region (175 nm) and is due to π π * transition which being symmetry allowed is intense. An isolated carbonyl group gives two absorption bands [190 (π π * ), 280 (n π * )]. The band at 190 nm is generally not observed as it lies below the cut off limits of the solvents. The other at 280 nm, though weak, lies much above the solvent cut off and is generally observed. However when the C=C and C= chromophores are conjugated (enones), both π π * and n π * transitions shift towards longer wavelength (red shift). In this case, an intense absorption band in the region nm is observed due to π π * transition and a weak absorption is observed at nm due to n π * transition. (Fig. 7.16) Woodward Fieser rules for calculation of λ max for π π * transition of α, β - unsaturated carbonyl compounds (enones) β α δ γ β α β C = C C = δ C = C C = C C = Base values 1. α, β- unsaturated acyclic or six membered ring ketone 215 nm 2. α, β- unsaturated five membered ring ketone 202 nm 3. α, β- unsaturated aldehydes 207 nm 18

19 4. α, β- unsaturated acids or esters 197 nm Increments (i) each alkyl group or ring residue α 10 nm β 12 nm γ or higher 18 nm (ii) each double bond extending conjugation 30 nm (iii) each exocyclic double bond 5 nm (iv) homoannular diene component 39 nm (i. e. two double bonds which are conjugated with C= lie in same ring) 5. Auxochrome α β γ δ - H (hydroxyl) 35 nm 30 nm 30 nm 50 nm - R (alkoxy) 35 nm 30 nm 17 nm 31 nm -C 6 nm 6 nm 6 nm 6 nm - Cl 15 nm 12 nm Br 25 nm 30 nm - - -NH2-95 nm Solvent correction Methanol 0 Cyclohexane +11 Chloroform +1 Dioxane +5 Ether +7 Water - 8 Hexane +11 Problems: Calculate λ max for the following compounds. CH 2 H 3 C 19

20 CH MeC H 3 C C H 3 Applications of UV spectroscopy Detection of functional groups UV spectroscopy has been used to detect the presence or absence of certain functional groups in organic molecules. For example, if a molecule doe not show absorption above 200 nm, it can be inferred that no conjugated chromophore, aldehyde, ketone, aromatic group is present. However an isolated alkene, alkyne may be present. Distinction between conjugated and non conjugated compounds can be made on the basis of observed λ max. Quantitative Analysis Due to high sensitivity UV Visible spectroscopy is a valuable analytical tool. By using Lambert Beer law, (A = εcl) which relates absorbance with concentration, quantitative analysis can be performed for compounds which show absorption in this region. Identification of geometrical isomers 20

21 Generally geometrical isomers absorbs at different wavelength. This fact is employed to distinguish between cis and trans alkene isomers. Trans isomers absorb at longer wavelength. λ max for trans stilbene is 294 nm and for cis stilbene it is 278 nm Detection of impurities Impurities, which absorb in UV region, can be detected from the UV spectra. A practical application is the detection of benzene even if present in low concentration in a sample of ethanol. Benzene absorbs at 254 nm where as ethanol does not absorbs above 200 nm. Similarly, impurity profiling of various drugs is done by using this technique. Detection of formation of charge transfer complexes Many electron rich organic molecules such as aromatic hydrocarbons, phenols, aromatic amines form charge transfer complexes with electron deficient molecule such as picric acid. During the formation of these complexes, the electrons are transferred from the donor to the acceptor resulting in strong charge transfer absorptions. Monitoring progress of a reaction UV spectra help in monitoring the course of a reaction by recording the spectrum of small amount of solution taken from reaction mixture at different intervals of the reaction. For example, reduction of a ketone to alcohol can be monitored by analyzing the n-π* transition of ketone. The absorption band reduces in intensity and confirms the completion of the reaction if disappears completely. 21