Ultraviolet Visible Molecular Absorption Spectroscopy old... · Ultraviolet Visible Molecular...
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Absorption Spectroscopy Absorption Spectroscopy Chapters 13 & 14 Skoog Holler Nieman 5th ed.
• The range of wavelengths for common UV/VIS is 180 nm < λ < 700 nm. The energy range can be described:
• The photon energy is typically described as kJ/mol for the UV/VIS region.
• The type of quantum transition is Bonding Electrons. There are four transitions.
• Limited to chromophores.• Transition Probability and cross sectional area of
absorbing species dictates magnitude of absorption.• Subject to solvent effects.
“Absorption measurements based on UV and VIS radiation find widespread application for the identification and determination of myriad organic and inorganic species. Molecular UV/VIS absorption methods are perhaps the most widely used of all quantitative analysis techniques in chemical and clinical laboratories throughout the world.”
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• E.g. A solution containing a UV absorbing chromophore (λmax 250 nm) is impinged with UV radiation (180 to 350 nm) of power PO (note: this is not a monochromatic light source) (UV/VIS-PDA)
• Some of PO is absorbed due to electron excitation within the chromophore, and is related to [analyte] and transition probability
• What emerges from the other side of the sample of a specific path length is the output power, P.
• The ratio of P to PO forms the basis of transmittance (T) and absorption (A).
• Beer’s Law
Beer’s LawSHN pp302-303
• The derivation of this law is well described.
TPPbcA o loglog −==⋅⋅= ε
(cm)length path cell theis (moles/L)ion concentrat analyte theis
)cm(Mt coefficien extinctionmolar theis
• The sample path length, the analyte concentration, its medium, its cross sectional area, and most importantly its extinction coefficient, dictate the resultant P relative to PO
The extinction coefficient
• The extinction coefficient is the term that is seldom thoroughly discussed in carrying out UV/VIS.
• Rearranging Beer’s Law to solve for ε indicates that it is inversely dependent on concentration and path length, and directly so on absorbance.
• The ε is actually a complex function which, for any particular peak in a spectrum, is dependent on the cross sectional area of the molecule (As in cm²) and its probability of transition (P);
• The transition probability, or oscillator strength, is actually a dimensionless value, which relates the integrated absorption coefficient AC for all the frequencies in the band.
s19 AP1078 ⋅⋅×= .ε
More on Extinction Coefficient
• Intense electronic transitions approach transition probabilities of unity, while significantly less intense transitions have transition probabilities significantly less than unity (orders of magnitude).
• The electronic transition from ground state to a higher energy state would yield a transition probability of unity for all possible transitions if the electron could vibrate in all three dimensions.
• The integrated absorption coefficient is hidden within the transition probability, but is quite a significant component. Asstated, the AC is the sum of all the intensities of all the transitions, so the greater it is, the greater is the transitionprobability.
• Calculation of the transition probability is possible using the transition dipole moment, µ. The transition dipole moment requires information regarding wavefunctions of the initial and final states of the molecule. It is therefore possible to calculate transition probability with a theoretical approach, provided thewavefunctions are known.
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Extinction Coefficient Ranges
• Typical capture cross section areas are (for organics) ~10–15cm².
• With transition probabilities ranging from: 0< P <1, we can realize extinction coefficients generally ranging from 10² to 105. Note that transition probabilities <0.01 are called “forbidden” transitions, in other words, not likely to occur relative to other more favorable routes of electronic excitation.
• In summary, – εmax ~104 to 105 yields strong absorption (intensity)– εmax ~103 yields medium absorption (intensity)– εmax <103 yields low absorption (intensity)
• Recall, absorption is the process of exciting an electron from one energy level to another, of which there are four possible transitions.
s19 AP1078 ⋅⋅×= .ε
σ, π, and lone pair (n) electrons
• The σ, π, and lone pair (n) electrons can all undergo excitation from ground state to excited energy states. Take a look at a depiction of the various electrons within a simple molecule, say formaldehyde.
• The four transitions are σ to σ*, π to π*, n to σ*, and n to π*. Recall that the promoted electrons are usually bonding electrons.
σ to σ* Transitions
• These are the bonding electrons. <150 nm
• The energies required to excite an electron from the ground state σ to excited state σ* is large -recall:
• Exciting an electron from σto σ* requires considerable energy, actually two to three times more energy compared to other transitions.
• Conventional UV spectroscopy cannot be used to detect these transitions.
– Absorption of radiation by glass and quartz
– The compounds need to be in a vacuum. (can you explain why?)
n to σ* Transitions
• These are the lone pair electrons or nonbonding electrons. 150 to 250nm.
• The transitions energies are lower than those of σto σ*.
• The extinction coefficients are small.
• Note that there is a hypsochromic shift upon exposure of these compounds to polar solvents. Why?
• The number of molecules with n to σ* transitions is small.
n to π* Transitions
• Unsaturated molecules that contain atoms such as oxygen, nitrogen, and sulfur exhibit a weak band in their absorption spectra which can be assigned to n to π*transitions.
• For aldehydes and ketones, this band arises from excitation of a non-bonding (lone-pair) electron on the oxygen atom into the antibonding π* orbital of the carbonyl moiety.
• In saturated aldehydes and ketones, the band due to n to π* transition generally occurs in the range between 270-300 nm, with reasonably high εmax.
• In unsaturated carbonyl compounds, the spectrum is usually a composite of the ethylenic moieties and those of the carbonyl function.
• If the carbonyl group is part of a conjugated system, the band due to nto π* transition generally appears in the range of 300-350 nm, and has an εmax substantially greater than that of the carbonyl alone (~10x greater).
...n to π* Transitions
• The addition of species such as OH and Cl, act as electron donors to the πsystem of the carbonyl group. This causes an increases in energy of all the πorbitals, bonding and anti-bonding, but doesn’t affect the energy of the n orbital. This causes an increases in the energy difference between the n and π* orbitals and thus we see an increase in the energy required to effect the transition, i.e., shift to shorter wavelength.
• The pH of the medium has a large effect on n to π* transitions. Consider the protonation of pyridine. This results in what appears to be a disappearance due to lone-pair bonding with proton. Under basic conditions (at a pH >> pKa), absorption at 300 nm is observed, which for interest is not present in benzene.
• Blue shift or Hypsochromic shift: Hydrogen bonding has significant impacts on n to π* transitions, that results in a blue shift, or a shift to shorter wavelengths. Consider a compound with a carbonyl group in hexane. The addition of a ethanol (through hydrogen bonding and with energy ~ equal to the hydrogen bond) will lower the energy level of the unexcited state of the n orbital, thus increasing the energy required to effect the transition. Wavelength shifts of 30 nm are not uncommon here.
• Increasing the polarity of the solvent has an effect on the n to π*, but not as significantly as does a change in polarity.
• εmax ranges from ~10 to 100 M-1cm-1
π to π* Transitions
• Characterized by the promotion of an electron from a bonding πorbital to an anti-bonding π* orbital.
• Selection rules, such as those based on symmetry concepts, determine whether a transition to a particular π* orbital is allowed or forbidden.
• There are several possibilities of π to π* transitions in molecules containing several πorbitals constituting the π-electron system. The extent to which transitions are allowed or forbidden varies, and spectra may exhibit bands with a range of intensities (e.g., ethylene and benzene).
• Ethylene: The allowed transition, Ag to B1utransition is intense (174 nm).
...π to π* Transitions
• Benzene: – A1g to E2u is an
allowed transition (~180 nm)(εmax~60,000)
– A1g to B1u (E2 band) (~204 nm) (εmax~8,000) and A1gto B2u (B band) (~256 nm)(εmax~200) are forbidden.
– Note that these bands are strongly affected by ring substitution.
...π to π* Transitions
• The effect of solvents on the position of bands due to πto π* transitions is usually opposite to that for n to π* transitions. A hydrolytic solvent as a medium for a molecule with potential π to π* transitions will cause a red shift or bathochromic shift in wavelength.
• Generally speaking, this is due to the fact the excited state in a π to π* transition is more stable in a hydrolytic solvent compared to a more hydrophobic solvent.
• An increase in the extent of conjugation on aromatic compounds causes the absorption bands to shift to longer wavelengths.
• For aromatic species, the effect of substituting proton in benzene with various species is quite marked.
• Generally, π to π* εmax ranges from ~1,000 to 10,000 M-
• Auxochromes:– Do not absorb in UV, but can shift chromophore
peaks to longer wavelengths (red shift) and can also increase their intensities.
– Note that both -OH and NH2 have an auxochromiceffect the B band of benzene.
– The auxochromic substituents have at least one pair of non-bonding electrons (lone pair) that can interact with the πelectrons of the ring (stabilizes the π* state).
Cut-off Points of Some Common Solvents
Absorption Characteristics of Some Common Chromophores
• There are a number of variables that influence absorbance:– nature of the solvent– pH of the solution– temperature (refractive index changes)– electrolyte concentration– interfering substances
Some UV Spectra
• Note that these are NOT “fingerprint” spectra.• There are huge solvent effects.
Electronic Energy Levels Including Transitions from
Vibrational and Rotational Levels
• We have to consider that each energy level has numerous vibrational energy levels, and those have numerous rotational energy levels.
• The electronic transition can therefore occur from energy levels above the pure electronic ground level energy level. This results in the broad distribution of absorption seen in UV/VIS.
–There are three basic processes by which a molecule can absorb radiation; all involve raising the molecule to a higher internal energy level, the increase in energy being equal to the energy of the absorbed radiation (hv). The three types of internal energy are quantized; that is, they exist at discrete levels.–First, the molecule rotates about various axes, the energy of rotation being at definite energy levels, so the molecule may absorb radiation and be raised to a higher rotational energy level, in a rotational transition. –Second, the atoms or groups of atoms within a molecule vibraterelative to each other, and the energy of this vibration occurs at definite quantized levels. The molecule may then absorb a discrete amount of energy and be raised to a higher vibrational energy level, in a vibrational transition. –Third, the electrons of a molecule may be raised to a higher level of electron energy, corresponding to an electronic transition. Since each of these internal energy transitions is quantized, they will occur only at definite wavelengths corresponding to an energy (hv) equal to the quantized jump in the internal energy. –There are, however, many different possible energy levels for each type of transition, and several wavelengths may be absorbed. The transitions can be illustrated by an energy level diagram. The relative energy levels of the three transition processes are in the order electronic > vibrational > rotational, each being about an order of magnitude different in its energy level. –Rotational transitions thus can take place at very low energies (long wavelengths, that is, the microwave or far-infrared region)–Vibrational transitions require higher energies in the near-infrared region–Electronic transitions require still higher energies (in the visible and ultraviolet regions).–Christain, G.D., Analytical Chemistry 5th Edition, 1994, pp. 402