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Electronic Spectroscopy of molecules. Regions of Electromagnetic Spectrum Radio-waves Region Microwaves Region Infra-red Region Visible Region Ultra-violet.

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Presentation on theme: "Electronic Spectroscopy of molecules. Regions of Electromagnetic Spectrum Radio-waves Region Microwaves Region Infra-red Region Visible Region Ultra-violet."— Presentation transcript:

1 Electronic Spectroscopy of molecules

2 Regions of Electromagnetic Spectrum Radio-waves Region Microwaves Region Infra-red Region Visible Region Ultra-violet Region X-ray Region  -ray Region Frequency (HZ) 10 6 - 10 10 10 10 - 10 12 10 12 - 10 14 10 14 - 10 15 10 15 - 10 16 10 16 - 10 18 10 18 - 10 20 Wavelength10m – 1 cm1 cm – 100µm100µm – 1µm700 – 400 nm400-10 nm10nm – 100pm 100pm – 1 pm NMR, ESRRotational Spectroscopy Vibrational spectroscopy Electronic Spectroscopy Electronic Spec. Energy0.001 – 10 J/mole Order of some 100 J/mole Some 10 4 J/mole Some 100 kJ/mole Some 100s kJ/mole 10 7 - 10 9 J/mole 10 9 - 10 11 J/mole Frequency ( ) Wavelength ( )

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

4 UVX-rays IR  -rays RadioMicrowave Visible UV Spectroscopy Longer Wavelength, Lower Energy

5  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 Spectroscopy Visible range UltraViolet range Longer Wavelength, Lower Energy 150 nm200 nm

6 Flame Test for Cations lithium sodium potassiumcopper 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 Flame Test 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. Light Photon

8 Emission of Energy (2 Possibilities) Continuous Energy LossQuantized Energy Loss or

9 Emission of Energy 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

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

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 = 1 n = 2 n = 3 n = 4 n = 5 n = 6 n = 7 n = 8

13 Line Emission Spectrum of Hydrogen Atoms H 2 Emission Spectrum

14 Line Spectra vs. Continuous Emission Spectra  The fact that the emission spectra of H 2 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 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. The Born-Oppenheimer Approximation: Different Types of Molecular Energy in Electronic Spectra A change in the total energy of a molecule may then by written,

17

18 Origin of Electronic Spectra

19 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. Origin of Electronic Spectra  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

20 Origin of Electronic Spectra

21 UV Spectrum of Isoprene

22 Types of Electrons in Molecules

23  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 ( ,  ), there is a corresponding anti-bonding orbital of symmetrically higher energy (  *,  * )  The lowest energy occupied orbitals are typically the  likewise, the corresponding anti-bonding   orbital is of the highest energy   -orbitals are of somewhat higher energy, and their complementary anti-bonding orbital somewhat lower in energy than  *.  Unshared pairs lie at the energy of the original atomic orbital, most often this energy is higher than  or  (since no bond is formed, there is no benefit in energy) Possible Electronic Transitions

24 Observed electronic transitions: graphical representation Energy     n Atomic orbital Molecular orbitals Occupied levels Unoccupied levels  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

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

26 MO’s Derived From the 2p Orbitals y y

27 Pi (  ) – 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

28 Observed electronic transitions: From the molecular orbital diagram, there are several possible electronic transitions that can occur, each of a different relative energy: Energy     n nnnn  alkanes carbonyls unsaturated cmpds. O, N, S, halogens carbonyls UV Spectroscopy

29 Possible Electronic Transitions

30 30 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: nnnn  alkanes carbonyls unsaturated cmpds. O, N, S, halogens carbonyls 150 nm 170 nm 180 nm √ - if conjugated! 190 nm 300 nm √ UV Spectroscopy

31 UV Spectrum of Isoprene

32 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

33 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

34 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  systems or carbonyls Common solvents and cutoffs: acetonitrile 190 chloroform240 cyclohexane195 1,4-dioxane215 95% ethanol205 n-hexane201 methanol205 isooctane195 water190 UV Spectroscopy

35 35 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 max max = 206 nm 252 317 376 UV Spectroscopy Wavelength (nm) Abs

36 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........ l = width of the cuvette A =  lc = log I 0 /I Where A is absorbance  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 Beer-Lambert Law:  = intensity of the incident light = intensity of the transmitted light Incident light Transmitted light A = log (Original intensity/ Intensity) % T = log (Intensity/ Original intensity) x 100

37  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 = log 10 (I 0 /I) or  log 10 (I/I 0 )  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,  this may vary by orders of magnitude… UV Spectroscopy A =  lc = log I 0 /I

38 The Spectrum These effects are combined into the Beer-Lambert Law:A =  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 10 4 -10 6 are termed high intensity absorptions values of 10 3 -10 4 are termed low intensity absorptions values of 0 to 10 3 are the absorptions of forbidden transitions A is unitless, so the units for  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 , and the y-axis is expressed as  directly or as the logarithm of  UV Spectroscopy

39 Observed electronic transitions: From the molecular orbital diagram, there are several possible electronic transitions that can occur, each of a different relative energy: Energy     n nnnn  alkanes carbonyls unsaturated cmpds. O, N, S, halogens carbonyls UV Spectroscopy: Electronic transitions

40 n   *,    * requires unsaturated functional groups most commonly used, energy good range for UV/Vis  ~ 200 - 700 nm n   * :  ~ 10-100    * :  ~ 1000 – 10,000 The valence electrons are the only ones whose energies permit them to be excited by near UV/visible radiation.  (bonding)  (bonding) n (non-bonding)  (anti-bonding)  (anti-bonding) Four types of transitions  * n  * n  *  *  * transition in vacuum UV ( ~ 150 nm) n  * saturated compounds with non- bonding electrons ~ 150-250 nm  ~ 100-3000 ( not strong)

41 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 NO 2 etc.  Structural or electronic changes in the chromophore can be quantified and used to predict shifts in the observed electronic transitions. UV Spectroscopy: Chromophores

42  Alkanes (CH 4, C 2 H 6 etc.) – only posses  -bonds and no lone pairs of electrons, so only the high energy    * transition is observed in the far UV (or vacuum UV), ~ 150 nm   UV Spectroscopy: Organic Chromophores

43  Alcohols, ethers, amines and sulfur compounds – in these compounds the n   * is the most often observed transition at shorter value (< 200 nm); like the alkane    * transition also possible. Note how this transition occurs from the HOMO to the LUMO  CN  CN n N sp 3 UV Spectroscopy: Organic Chromophores

44 n  * transition lower in energy than σ  * max  max H 2 O1671480 CH 3 OH184150 CH 3 Cl173200 CH 3 I258365 (CH 3 ) 2 S229140 (CH 3 ) 2 O1842520 CH 3 NH 2 215600 (CH 3 ) 3 N227900 UV Spectroscopy: Chromophores Alcohols, ethers, amines and sulfur compounds Explain why max and the corresponding  max is different. n  * transition - - between 150 and 250 nm. n  * transition

45  Alkenes and Alkynes – in the case of isolated examples of these compounds the    * is observed at 175 and 170 nm, respectively Even though this transition is of lower energy than    *, it is still in the far UV – however, the transition energy is sensitive to substitution   UV Spectroscopy: Organic Chromophores ~ 170 - 190 nm

46 Example: ethylene absorbs at max = 165 nm  = 10,000 (intense band)  = hv =hc/   hv       UV Spectroscopy: Organic Chromophores Alkenes

47 Carbonyls – unsaturated systems incorporating N or O can undergo n   * transitions (~285 nm) in addition to    * Despite the fact this transition is forbidden by the selection rules (  = 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    * transition in the vacuum UV (188 nm,  = 900); sensitive to substitution effects UV Spectroscopy: Organic Chromophores n  *n  *

48 Carbonyls – n   * transitions (~285 nm);    * (188 nm)   n σ  σ * transitions omitted for clarity UV Spectroscopy: Organic Chromophores

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

50 n  * and  * Transitions Most UV/vis spectra involve these transitions.  * are generally more intense than n  * max  max type C 6 H 13 CH=CH 2 17713000  * C 5 H 11 C  C–CH 3 17810000  * O CH 3 CCH 3 1861000 n  * O CH 3 COH204 41 n  * CH 3 NO 2 280 22 n  * CH 3 N=NCH 3 339 5 n  * UV Spectroscopy: Chromophores

51 ChromophoreExampleSolvent max (nm)  max Type of transition Alkenen-Heptane17713,000 ** Alkyne C 5 H 11 C  C-CH 3 n-Heptane178 196 225 10,000 2,000 160 *__*__ Carbonyln-Hexane 186 280 180 293 1,000 16 Large 12 n*n*n*n*n*n*n*n* CarboxylEthanol20441 n*n* AmidoWater21460 n*n* AzoEthanol3395 n*n* Nitro CH 3 NO 2 Isooctane28022 n*n* Nitroso C 4 H 9 NO Ethyl ether300 665 100 20 _n*_n* Nitrate C 2 H 5 ONO 2 Dioxane27012 n*n* Absorption Characteristics of Some Common Chromophores

52 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 -CH 3, Et etc), hydroxyl (-OH), alkoxy (-OR) and amino groups (-NR 2 ) and the halogens (such as X = Cl, I etc.) UV Spectroscopy: Chromophores

53 In General – Substituents may have any of four effects on a chromophore i.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 200 nm700 nm  Hypochromic Hypsochromic Hyperchromic Bathochromic UV Spectroscopy: Substituent Effects

54 Conjugation – most efficient means of bringing about a bathochromic and hyperchromic shift of an unsaturated chromophore: max nm  17515,000 21721,000 25835,000 n   *280 27    *213 7,100 465125,000 n   *280 12    *189 900 UV Spectroscopy: Substituent Effects


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