UV Spectroscopy
Electronic Spectroscopy of molecules Regions of Electromagnetic Spectrum Radio-waves Region Microwaves Region Infra-red Region Visible Region Ultra-violet Region X-ray Region -ray Region Frequency (HZ) 106 - 1010 1010 - 1012 1012 - 1014 1014 - 1015 1015 - 1016 1016 - 1018 1018- 1020 Wavelength 10m – 1 cm 1 cm – 100µm 100µm – 1µm 700 – 400 nm 400-10 nm 10nm – 100pm 100pm – 1 pm NMR, ESR Rotational Spectroscopy Vibrational spectroscopy Electronic Spectroscopy Electronic Spec. 0.001 – 10 J/mole Order of some 100 J/mole Some 104 J/mole Some 100 kJ/mole Some 100s kJ/mole 107- 109 J/mole 109- 1011 J/mole Energy Frequency () Wavelength () Electromagnetic Radiation Energy of light Frequency of light E = h Higher frequency () -- Higher Energy -- Lower wavelength where h = Planck’s constant = 6.624 x 10-34 Joules sec = frequency of electromagnetic radiation in cycle per sec = c/ where c = velocity of light; = wavelength of electromagnetic radiation Therefore, E = hc/ But 1/ = = wave number in cm-1 Thus, E = hc UV Spectroscopy -rays X-rays UV IR Microwave Radio Visible Longer Wavelength, Lower Energy UltraViolet Spectroscopy Also known as electronic spectroscopy Involves transition of electrons within a molecule or ion from a lower to higher electronic energy level or vice versa by absorption or emission of radiation falling in the uv-visible range, Visible range is 400-800 nm Near uv is 200-400 nm Far uv is 150-200 nms UltraViolet range Visible range 150 nm 200 nm Longer Wavelength, Lower Energy Flame Test for Cations lithium sodium potassium copper A flame test is an analytic procedure used in chemistry to detect the presence of certain elements, primarily metal ions, based on each element's characteristic emission spectrum. The color of flames in general also depends on temperature. Flame Test Light Photon 1. Electron absorbs energy from the flame goes to a higher energy state. 2. Electron goes back down to lower energy state and releases the energy it absorbed as light. Emission of Energy (2 Possibilities) or Continuous Energy Loss Quantized Energy Loss Emission of Energy Continuous Energy Loss Quantized Energy Loss • Any and all energy values possible on way down • Implies electron can be anywhere about nucleus of atom • Only certain, restricted, quantized energy values possible on way down • Implies an electron is restricted to quantized energy levels Continuous emission spectra Line spectra Emission Spectrum Continuous Emission Spectrum Line Emission Spectrum (Quantized Energy Loss) Atomic Spectra of Hydrogen Atom http://hyperphysics.phy-astr.gsu.edu/hbase/hyde.html Atomic Spectra of Hydrogen Atom by Bohr´s Theory: n=8 n=7 n=6 n=5 n=4 n=3 n=2 n=1 H2 Emission Spectrum Line Emission Spectrum of Hydrogen Atoms Line Spectra vs. Continuous Emission Spectra The fact that the emission spectra of H2 gas and other molecules is a line rather than continuous emission spectra tells us that electrons are in quantized energy levels rather than anywhere about nucleus of atom. Regions of Electromagnetic Spectrum Different Types of Molecular Energy in Electronic Spectra The Born-Oppenheimer Approximation: A change in the total energy of a molecule may then by written, Pure rotational spectra: Permanent electric dipole moment – Fine Structure IR or Vibrational Spectra: Change of dipole during motion Electronic Spectra: Changes in electron distribution in a molecule are always accompanied by a dipole change. ALL MOLECULES DO GIVE AN ELECTRONIC SPECTRUM AND SHOW VIBRATIONAL AND ROTATIONAL STRUCTURE IN THEIR SPECTRA FROM WHICH ROTATIONAL CONSTANTS AND BOND VIBRATION FREQUENCIES MAY BE DERIVED. Origin of Electronic Spectra Origin of Electronic Spectra In the ground state electrons are paired If transition of electron from ground state to excited state takes place in such a way that spins of electrons are paired, it is known as excited singlet state. If electrons have parallel spins, it is known as excited triplet state. Excitation of uv light results in excitation of electron from singlet ground state to singlet excited state Transition from singlet ground state to excited triplet state is forbidden due to symmetry consideration Origin of Electronic Spectra UV Spectrum of Isoprene Types of Electrons in Molecules Possible Electronic Transitions The lowest energy transition (and most often obs. by UV) is typically that of an electron in the Highest Occupied Molecular Orbital (HOMO) to the Lowest Unoccupied Molecular Orbital (LUMO). For any bond (pair of electrons) in a molecule, the molecular orbitals are a mixture of the two contributing atomic orbitals; for every bonding orbital “created” from this mixing (s, p), there is a corresponding anti-bonding orbital of symmetrically higher energy (s*, p*) The lowest energy occupied orbitals are typically the s; likewise, the corresponding anti-bonding s* orbital is of the highest energy p-orbitals are of somewhat higher energy, and their complementary anti-bonding orbital somewhat lower in energy than s*. Unshared pairs lie at the energy of the original atomic orbital, most often this energy is higher than p or s (since no bond is formed, there is no benefit in energy) Observed electronic transitions: graphical representation s* Unoccupied levels p* Energy Atomic orbital n Atomic orbital Occupied levels p s Molecular orbitals The difference in energy between molecular bonding, non-bonding and anti-bonding orbitals ranges from 125-650 kJ/mole This energy corresponds to EM radiation in the ultraviolet (UV) region, 100-350 nm, and visible (VIS) regions 350-700 nm of the spectrum UV Spectroscopy - Single bonds are usually too high in excitation energy for most instruments (185 nm) -- vacuum UV. Types of electron transitions: Sigma (s) – single bond electron i) s, p, n electrons A MO A MO’s Derived From the 2p Orbitals y y UV Spectroscopy Pi (p) – double bond electron Low energy bonding orbital (π) High energy anti-bonding orbital (π*) Non-bonding electrons (n): don’t take part in any bonds -- neutral energy level. Example: Formaldehyde UV Spectroscopy Observed electronic transitions: From the molecular orbital diagram, there are several possible electronic transitions that can occur, each of a different relative energy: s* p* Energy n p s s s* alkanes s p* carbonyls p p* unsaturated cmpds. n s* O, N, S, halogens n p* carbonyls Possible Electronic Transitions UV Spectroscopy Observed electronic transitions: • Although the UV spectrum extends below 100 nm (high energy), oxygen in the atmosphere is not transparent below 200 nm • Special equipment to study vacuum or far UV is required • Routine organic UV spectra are typically collected from 200-700 nm • This limits the transitions that can be observed: s s* alkanes 150 nm s p* carbonyls 170 nm p p* unsaturated cmpds. 180 nm n s* O, N, S, halogens 190 nm n p* carbonyls 300 nm √ - if conjugated! √ 30 UV Spectrum of Isoprene UV Spectroscopy Selection Rules Not all transitions that are possible are observed For an electron to transition, certain quantum mechanical constraints apply – these are called “selection rules” For example, an electron cannot change its spin quantum number during a transition – these are “forbidden” Other examples include: • the number of electrons that can be excited at one time • symmetry properties of the molecule • symmetry of the electronic states To further complicate matters, “forbidden” transitions are sometimes observed (albeit at low intensity) due to other factors..... UV Spectroscopy Instrumentation – Sample Handling In general, UV spectra are recorded solution-phase Cells can be made of plastic, glass or quartz Only quartz is transparent in the full 200-700 nm range; plastic and glass are only suitable for visible spectra Concentration is empirically determined A typical sample cell (commonly called a cuvet): UV Spectroscopy Instrumentation – Sample Handling Solvents must be transparent in the region to be observed; the wavelength where a solvent is no longer transparent is referred to as the cutoff Since spectra are only obtained up to 200 nm, solvents typically only need to lack conjugated p systems or carbonyls Common solvents and cutoffs: acetonitrile chloroform cyclohexane 1,4-dioxane 95% ethanol n-hexane methanol isooctane water 190 240 195 215 205 201 205 195 190 UV Spectroscopy The Spectrum The x-axis of the spectrum is in wavelength; 200-350 nm for UV, 200-700 for UVVIS determinations Due to the lack of any fine structure, spectra are rarely shown in their raw form, rather, the peak maxima are simply reported as a numerical list of “lamba max” values or max NH2 max = Abs O Wavelength (nm) O 206 nm 252 317 376 35 Beer-Lambert Law: When a beam of monochromatic radiation is passed through a solution of an absorbing medium, the rate of decrease of intensity of radiation with thickness of the absorbing medium is directly proportional to the intensity of incident radiation as well as the concentration of the solution........ A = elc = log I0/I Where A is absorbance e is the molar absorbtivity with units of L mol-1 cm-1 l is the path length of the sample (typically in cm). c is the concentration of the compound in solution, expressed in mol L-1 = intensity of the incident light Transmitted light Incident light = intensity of the transmitted light l = width of the cuvette e A = log (Original intensity/ Intensity) % T = log (Intensity/ Original intensity) x 100 UV Spectroscopy The Spectrum The y-axis of the spectrum is in absorbance, A From the spectrometers point of view, absorbance is the inverse of transmittance: A = log10 (I0/I) or log10 (I/I0) From an experimental point of view, three other considerations must be made: a longer path length (l ) through the sample will cause more UV light to be absorbed – linear effect the greater the concentration (c) of the sample, the more UV light will be absorbed – linear effect some electronic transitions are more effective at the absorption of photon than others – molar absorptivity, e this may vary by orders of magnitude… A = elc = log I0/I e UV Spectroscopy The Spectrum These effects are combined into the Beer-Lambert Law: A=ecl for most UV spectrometers, l would remain constant (standard cells are typically 1 cm in path length) concentration is typically varied depending on the strength of absorption observed or expected – typically dilute – sub .001 M molar absorptivities vary by orders of magnitude: • values of 104-106 are termed high intensity absorptions • values of 103-104 are termed low intensity absorptions • values of 0 to 103 are the absorptions of forbidden transitions A is unitless, so the units for e are cm-1 · M-1 and are rarely expressed Since path length and concentration effects can be easily factored out, absorbance simply becomes proportional to e, and the y-axis is expressed as e directly or as the logarithm of e. UV Spectroscopy: Electronic transitions Observed electronic transitions: From the molecular orbital diagram, there are several possible electronic transitions that can occur, each of a different relative energy: s* p* Energy n p s s s* alkanes s p* carbonyls p p* unsaturated cmpds. n s* O, N, S, halogens n p* carbonyls The valence electrons are the only ones whose energies permit them to be excited by near UV/visible radiation. s* (anti-bonding) p* (anti-bonding) Four types of transitions ss* pp* n (non-bonding) ns* np* p (bonding) s (bonding) ss* transition in vacuum UV ( ~ 150 nm) ns* saturated compounds with nonbonding electrons ~ 150-250 nm e ~ 100-3000 ( not strong) n p*, p p* requires unsaturated functional groups most commonly used, energy good range for UV/Vis ~ 200 - 700 nm n p* : e ~ 10-100 p p* : e ~ 1000 – 10,000 UV Spectroscopy: Chromophores Definition Remember the electrons present in organic molecules are involved in covalent bonds or lone pairs of electrons on hetero-atoms such as O or N Since similar functional groups will have electrons capable of discrete classes of transitions, the characteristic energy of these energies is more representative of the functional group than the electrons themselves. A functional group capable of having characteristic electronic transitions is called a chromophore (color loving). A Chromophore is a covalently unsaturated group responsible for electronic absorption e.g C=C, C=0, and NO2 etc. Structural or electronic changes in the chromophore can be quantified and used to predict shifts in the observed electronic transitions. UV Spectroscopy: Organic Chromophores Alkanes (CH4, C2H6 etc.) – only posses s-bonds and no lone pairs of electrons, so only the high energy s s* transition is observed in the far UV (or vacuum UV), ~ 150 nm C s* s C C C UV Spectroscopy: Organic Chromophores Alcohols, ethers, amines and sulfur compounds – in these compounds the n s* is the most often observed transition at shorter value (< 200 nm); like the alkane s s* transition also possible. Note how this transition occurs from the HOMO to the LUMO s*CN C N nN sp C 3 C N N anitbonding orbital sCN C N UV Spectroscopy: Chromophores Alcohols, ethers, amines and sulfur compounds ns* transition lower in energy than σs* ns* transition - - between 150 and 250 nm. max emax H2O 167 1480 CH3OH 184 150 CH3Cl 173 200 CH3I 258 365 (CH3)2S 229 140 (CH3)2O 184 2520 CH3NH2 215 600 (CH3)3N 227 900 Explain why max and the corresponding emax is different. UV Spectroscopy: Organic Chromophores Alkenes and Alkynes – in the case of isolated examples of these compounds the p p* is observed at 175 and 170 nm, respectively Even though this transition is of lower energy than s s*, it is still in the far UV – however, the transition energy is sensitive to substitution p* ~ 170 - 190 nm p UV Spectroscopy: Organic Chromophores Alkenes C C s* p* s* p* = hv =hc/ hv p p s s p p* Example: ethylene absorbs at max = 165 nm e= 10,000 (intense band) UV Spectroscopy: Organic Chromophores C O n p* Carbonyls – unsaturated systems incorporating N or O can undergo n p* transitions (~285 nm) in addition to p p* Despite the fact this transition is forbidden by the selection rules (e = 15), it is the most often observed and studied transition for carbonyls This transition is also sensitive to substituents on the carbonyl Similar to alkenes and alkynes, non-substituted carbonyls undergo the p p* transition in the vacuum UV (188 nm, e = 900); sensitive to substitution effects UV Spectroscopy: Organic Chromophores C O Carbonyls – n p* transitions (~285 nm); p p* (188 nm) O p* n p σ σ* transitions omitted for clarity C O O UV Spectroscopy: Organic Chromophores s* p* C O n p* n s* p* hv p n p s s The np* transition is at even longer wavelengths (low energy transition) but is not as strong as pp* transitions. It is said to be “forbidden.” Example: Acetone: ns* max = 188 nm ; e= 1860 (intense band) np* max = 279 nm ; e= 15 UV Spectroscopy: Chromophores np* and pp* Transitions Most UV/vis spectra involve these transitions. pp* are generally more intense than np* max emax C6H13CH=CH2 177 13000 pp* C5H11CC–CH3 178 10000 pp* 186 1000 ns* CH3COH 204 41 np* CH3NO2 280 22 np* CH3N=NCH3 339 5 np* type O CH3CCH3 O Absorption Characteristics of Some Common Chromophores Chromophore Example Alkene C6H13HC Solvent CH2 Alkyne O emax Type of transition n-Heptane 177 13,000 pp* n-Heptane 178 196 225 10,000 2,000 160 pp* _ _ n-Hexane 186 280 1,000 16 ns* np* n-Hexane 180 293 Large 12 Ethanol 204 41 np* Water 214 60 np* Ethanol 339 5 np* C5H11CC-CH3 Carbonyl max (nm) CH3CCH3 O CH3CH Carboxyl Amido O CH3COH O ns* np* CH3CNH2 Azo H3CN NCH3 Nitro CH3NO2 Isooctane 280 22 np* Nitroso C4H9NO Ethyl ether 300 665 100 20 _ np* 270 12 np* Nitrate C2H5ONO2 Dioxane UV Spectroscopy: Chromophores Substituent Effects The attachment of substituent groups (other than H) can shift the energy of the transition Auxoxhromes: Substituents that increase the intensity and often wavelength of an absorption are called auxochromes. An auxochrome represents a saturated group, which when attached to a chromophore changes both the intensity as well as the wavelength of the absorption maximum. Common auxochromes include alkyl (such as -CH3, Et etc), hydroxyl (-OH), alkoxy (-OR) and amino groups (-NR2) and the halogens (such as X = Cl, I etc.) UV Spectroscopy: Substituent Effects In General – Substituents may have any of four effects on a chromophore Bathochromic shift (red shift) – a shift to longer ; lower energy ii. Hypsochromic shift (blue shift) – shift to shorter ; higher energy iii. Hyperchromic effect – an increase in intensity iv. Hypochromic effect – a decrease in intensity e Hypsochromic Bathochromic Hypochromic 200 nm Hyperchromic i. 700 nm UV Spectroscopy: Substituent Effects Conjugation – most efficient means of bringing about a bathochromic and hyperchromic shift of an unsaturated chromophore: H2C max nm CH2 e 175 15,000 217 21,000 258 35,000 465 125,000 -carotene O n p* 280 12 900 n p* 280 27 7,100 p p* 189 O p p* 213
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