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Electronic Spectroscopy of molecules

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1 Electronic Spectroscopy of molecules

2 Regions of Electromagnetic Spectrum
Radio-waves Region Microwaves Infra-red Visible Ultra-violet X-ray Region -ray Region Frequency (HZ) Wavelength 10m – 1 cm 1 cm – 100µm 100µm – 1µm 700 – 400 nm nm 10nm – 100pm 100pm – 1 pm NMR, ESR Rotational Spectroscopy Vibrational spectroscopy Electronic Spec. Energy 0.001 – 10 J/mole Order of some 100 J/mole Some 104 J/mole Some 100 kJ/mole Some 100s J/mole Frequency () Wavelength ()

3 Electromagnetic Radiation
Energy of light Frequency of light E = h Higher frequency () -- Higher Energy -- Lower wavelength where h = Planck’s constant = x 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

4 UV Spectroscopy IR g-rays UV X-rays Radio Microwave 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 nm Near uv is nm Far uv is nms UltraViolet range Visible range 150 nm 200 nm Longer Wavelength, Lower Energy

6 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.

7 Flame Test Electron absorbs energy from the flame goes to a higher energy state. Light Photon 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 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 Line Emission Spectrum (Quantized Energy Loss)
Continuous Emission Spectrum Line Emission Spectrum (Quantized Energy Loss)

11 Atomic Spectra of Hydrogen Atom

12 Atomic Spectra of Hydrogen Atom by Bohr´s Theory:

13 H2 Emission Spectrum Line Emission Spectrum of Hydrogen Atoms

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.


18 Origin of Electronic Spectra

19 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

20 Origin of Electronic Spectra

21 UV Spectrum of Isoprene

22 Types of Electrons in Molecules

23 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)

24 Observed electronic transitions: graphical representation
Unoccupied levels p* Energy Atomic orbital Atomic orbital n Occupied levels p s Molecular orbitals The difference in energy between molecular bonding, non-bonding and anti-bonding orbitals ranges from kJ/mole This energy corresponds to EM radiation in the ultraviolet (UV) region, nm, and visible (VIS) regions nm of the spectrum

25 Sigma (s) – single bond electron
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

26 MO’s Derived From the 2p Orbitals

27 Pi (p) – double bond electron
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

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

29 Possible Electronic Transitions

30 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 nm This limits the transitions that can be observed: s p n s* p* alkanes carbonyls unsaturated cmpds. O, N, S, halogens 150 nm 170 nm 180 nm √ - if conjugated! 190 nm 300 nm √

31 UV Spectrum of Isoprene

32 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.....

33 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 nm range; plastic and glass are only suitable for visible spectra Concentration is empirically determined A typical sample cell (commonly called a cuvet):

34 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

35 UV Spectroscopy The Spectrum
The x-axis of the spectrum is in wavelength; nm for UV, 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)

36 A = elc = log I0/I Beer-Lambert Law: l = width of the cuvette
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 e = intensity of the incident light Transmitted light = intensity of the transmitted light Incident light l = width of the cuvette A = log (Original intensity/ Intensity) % T = log (Intensity/ Original intensity) x 100

37 A = elc = log I0/I 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… e A = elc = log I0/I

38 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 are termed high intensity absorptions values of 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.

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

40 s* (anti-bonding) p* (anti-bonding) * * n (non-bonding) 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) * n* p* (anti-bonding) * n* n (non-bonding) p (bonding) s (bonding) * transition in vacuum UV ( ~ 150 nm) n* saturated compounds with non- bonding electrons  ~ nm e ~ ( not strong) n  p*, p  p* requires unsaturated functional groups most commonly used, energy good range for UV/Vis  ~ nm n  p* : e ~ p  p* : e ~ 1000 – 10,000

41 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.

42 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

43 UV Spectroscopy: Organic Chromophores
Alcohols, ethers, amines and sulfur compounds – in these compounds the n  s* is the most often observed transition at shorter l value (< 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

44 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 CH3OH CH3Cl CH3I (CH3)2S (CH3)2O CH3NH (CH3)3N n* transition Explain why max and the corresponding max is different.

45 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*  ~ nm p

46 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)

47 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

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

49 UV Spectroscopy: Organic Chromophores
hv n 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

50 UV Spectroscopy: Chromophores
n* and * Transitions Most UV/vis spectra involve these transitions. * are generally more intense than n* max max type C6H13CH=CH * C5H11CC–CH * O CH3CCH n* CH3COH n* CH3NO n* CH3N=NCH n*

51 Absorption Characteristics of Some Common Chromophores
Example Solvent lmax (nm) emax Type of transition Alkene n-Heptane 177 13,000 pp* Alkyne C5H11CC-CH3 178 196 225 10,000 2,000 160 _ Carbonyl n-Hexane 186 280 180 293 1,000 16 Large 12 ns* np* Carboxyl Ethanol 204 41 Amido Water 214 60 Azo 339 5 Nitro CH3NO2 Isooctane 22 Nitroso C4H9NO Ethyl ether 300 665 100 20 Nitrate C2H5ONO2 Dioxane 270

52 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.)

53 UV Spectroscopy: Substituent Effects
In General – Substituents may have any of four effects on a chromophore Bathochromic shift (red shift) – a shift to longer l; lower energy Hypsochromic shift (blue shift) – shift to shorter l; higher energy Hyperchromic effect – an increase in intensity Hypochromic effect – a decrease in intensity Hyperchromic e Hypsochromic Bathochromic Hypochromic 200 nm 700 nm

54 UV Spectroscopy: Substituent Effects
Conjugation – most efficient means of bringing about a bathochromic and hyperchromic shift of an unsaturated chromophore: lmax nm e 175 15,000 217 21,000 258 35,000 ,000 n  p* p  p* n  p* p  p* ,100

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