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Electronic Spectroscopy of molecules

Electronic Spectroscopy of molecules. Regions of Electromagnetic Spectrum. Frequency (  ). Wavelength ( ). Electromagnetic Radiation. Energy of light. Frequency of light. E = h . where h = Planck’s constant = 6.624 x 10 -34 Joules sec

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Electronic Spectroscopy of molecules

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  1. Electronic Spectroscopy of molecules

  2. Regions of Electromagnetic Spectrum Frequency () Wavelength ()

  3. Electromagnetic Radiation Energy of light Frequency of light E = h 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 Higher frequency () -- Higher Energy -- Lower wavelength  

  4. UV Spectroscopy g-rays X-rays UV IR Microwave Radio Visible Longer Wavelength, Lower Energy

  5. 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 Visible range UltraViolet range 150 nm 200 nm Longer Wavelength, Lower Energy

  6. Flame Test for Cations sodium lithium 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.

  7. Light Photon Flame Test • 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.

  8. Emission of Energy(2 Possibilities) or Continuous Energy Loss Quantized Energy Loss

  9. Continuous Energy Loss Any and all energy values possible on way down Implies electron can be anywhere about nucleus of atom Continuous emission spectra Quantized Energy Loss Only certain, restricted, quantized energy values possible on way down Implies an electron is restricted to quantized energy levels Line spectra Emission of Energy

  10. Emission Spectrum Line Emission Spectrum (Quantized Energy Loss) Continuous EmissionSpectrum

  11. Atomic Spectra of Hydrogen Atom http://hyperphysics.phy-astr.gsu.edu/hbase/hyde.html

  12. 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

  13. Line Emission Spectrum of Hydrogen Atoms H2 Emission Spectrum

  14. 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.

  15. Regions of Electromagnetic Spectrum

  16. 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.

  17. Origin of Electronic Spectra

  18. 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

  19. Origin of Electronic Spectra

  20. UV Spectrum of Isoprene

  21. Types of Electrons in Molecules

  22. 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)

  23. Observed electronic transitions: graphical representation s* Unoccupied levels p* Atomic orbital Atomic orbital Energy n 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

  24. UV Spectroscopy - Single bonds are usually too high in excitation energy for most instruments (185 nm) -- vacuum UV. Types of electron transitions: i) s, p, n electrons Sigma (s) – single bond electron A MO A

  25. MO’s Derived From the 2p Orbitals y y

  26. UV Spectroscopy Pi (p) – double bond electron High energy anti-bonding orbital (π*) Low energy bonding orbital (π) Non-bonding electrons (n): don’t take part in any bonds -- neutral energy level. Example: Formaldehyde

  27. 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* alkanes carbonyls unsaturated cmpds. O, N, S, halogens carbonyls s s p n n s* p* p* s* p* p* Energy n p s

  28. Possible Electronic Transitions

  29. 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: alkanes carbonyls unsaturated cmpds. O, N, S, halogens carbonyls 150 nm 170 nm 180 nm √ - if conjugated! 190 nm 300 nm √ s s p n n s* p* p* s* p*

  30. UV Spectrum of Isoprene

  31. 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.....

  32. 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):

  33. 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 190 • chloroform 240 • cyclohexane 195 • 1,4-dioxane 215 • 95% ethanol 205 • n-hexane 201 • methanol 205 • isooctane 195 • water 190

  34. UV Spectroscopy • The Spectrum • The x-axis of the spectrum is in wavelength; 200-350 nm for UV, 200-700 for UV-VIS 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 lmax lmax = 206 nm 252 317 376 Abs Wavelength (nm)

  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........ 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 A = elc = log I0/I = intensity of the incident light = intensity of the transmitted light Transmitted light Incident light l = width of the cuvette A = log (Original intensity/ Intensity) e % T = log (Intensity/ Original intensity) x 100

  36. 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

  37. UV Spectroscopy • The Spectrum • These effects are combined into the Beer-Lambert Law: A = e c l • 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.

  38. 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* alkanes carbonyls unsaturated cmpds. O, N, S, halogens carbonyls s s p n n s* p* p* s* p* p* Energy n p s

  39. * * n* n* The valence electrons are the only ones whose energies permit them to be excited by near UV/visible radiation. Four types of transitions s* (anti-bonding) p* (anti-bonding) n (non-bonding) p (bonding) s (bonding) * transition in vacuum UV ( ~ 150 nm) n* saturated compounds with non- bonding 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

  40. 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.

  41. 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 s* s

  42. UV Spectroscopy: Organic Chromophores • Alcohols, ethers, amines and sulfur compounds– in these compounds • the n s* is the most often observed transition at shorter lvalue (< 200 nm); • like the alkane s  s* transition also possible. • Note how this transition occurs from the HOMO to the LUMO s*CN nN sp3 sCN

  43. UV Spectroscopy: Chromophores Alcohols, ethers, amines and sulfur compounds n* transition lower in energy than σ* n* transition - -  between 150 and 250 nm. max max 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 n* transition Explain why max and the corresponding max is different.

  44. 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

  45. UV Spectroscopy: Organic Chromophores Alkenes s* s* p* p* = hv =hc/ hv p p s s Example: ethylene absorbs at max = 165 nm = 10,000(intense band)

  46. UV Spectroscopy: Organic Chromophores 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

  47. UV Spectroscopy: Organic Chromophores • Carbonyls – n p* transitions (~285 nm); p  p* (188 nm) p* n p σ σ* transitions omitted for clarity

  48. UV Spectroscopy: Organic Chromophores s* s* p* p* n n hv p p s s The np* transition is at even longer wavelengths (low energy transition) but is not as strong as pp* transitions. It is said to be “forbidden.” Example: Acetone: n max = 188 nm ; = 1860 (intense band)n max = 279 nm ; = 15

  49. UV Spectroscopy: Chromophores n* and * Transitions Most UV/vis spectra involve these transitions. * are generally more intense than n* max max type C6H13CH=CH2 177 13000 * C5H11CC–CH3 178 10000 * O CH3CCH3 186 1000 n* O CH3COH 204 41 n* CH3NO2 280 22 n* CH3N=NCH3 339 5 n*

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