2 Regions of Electromagnetic Spectrum Radio-wavesRegionMicrowavesInfra-redVisibleUltra-violetX-ray Region-ray RegionFrequency (HZ)Wavelength10m – 1 cm1 cm – 100µm100µm – 1µm700 – 400 nmnm10nm – 100pm100pm –1 pmNMR, ESRRotationalSpectroscopyVibrational spectroscopyElectronicSpec.Energy0.001 – 10 J/moleOrder of some 100 J/moleSome 104 J/moleSome 100kJ/moleSome 100sJ/moleFrequency ()Wavelength ()
3 Electromagnetic Radiation Energy of lightFrequency of lightE = hHigher frequency () -- Higher Energy -- Lower wavelengthwhere h = Planck’s constant = x Joules sec = frequency of electromagnetic radiation in cycle per sec = c/where c = velocity of light; = wavelength of electromagnetic radiationTherefore, E = hc/But 1/ = = wave number in cm-1Thus, 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 spectroscopyInvolves 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 nmNear uv is nmFar uv is nmsUltraViolet rangeVisible range150 nm200 nmLonger 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 TestElectron absorbsenergy from the flamegoes to a higher energystate.LightPhoton2. Electron goes back down to lower energy state and releases the energy it absorbed as light.
8 Emission of Energy (2 Possibilities) orContinuous Energy LossQuantized Energy Loss
9 Emission of Energy Continuous Energy Loss Any and all energy values possible on way downImplies electron can be anywhere about nucleus of atomContinuous emission spectraQuantized Energy LossOnly certain, restricted, quantized energy values possible on way downImplies an electron is restricted to quantized energy levelsLine spectra
10 Line Emission Spectrum (Quantized Energy Loss) Continuous Emission SpectrumLine Emission Spectrum (Quantized Energy Loss)
12 Atomic Spectra of Hydrogen Atom by Bohr´s Theory:
13 H2 Emission SpectrumLine 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.
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 StructureIR or Vibrational Spectra: Change of dipole during motionElectronic Spectra: Changes in electron distribution in a molecule are always accompaniedby 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.
19 Origin of Electronic Spectra In the ground state electrons are pairedIf 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 stateTransition from singlet ground state to excited triplet state is forbidden due to symmetry consideration
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 energyp-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 levelsp*EnergyAtomic orbitalAtomic orbitalnOccupied levelspsMolecular orbitalsThe difference in energy between molecular bonding, non-bonding and anti-bonding orbitals ranges from kJ/moleThis 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 electronsSigma (s) – single bond electronAMOA
27 Pi (p) – double bond electron UV SpectroscopyPi (p) – double bond electronHigh 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*spns*p*alkanescarbonylsunsaturated cmpds.O, N, S, halogensp*Energynps
30 UV SpectroscopyObserved electronic transitions:Although the UV spectrum extends below 100 nm (high energy), oxygen in the atmosphere is not transparent below 200 nmSpecial equipment to study vacuum or far UV is requiredRoutine organic UV spectra are typically collected from nmThis limits the transitions that can be observed:spns*p*alkanescarbonylsunsaturated cmpds.O, N, S, halogens150 nm170 nm180 nm √ - if conjugated!190 nm300 nm √
32 UV SpectroscopySelection RulesNot all transitions that are possible are observedFor 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 timesymmetry properties of the moleculesymmetry of the electronic statesTo further complicate matters, “forbidden” transitions are sometimes observed (albeit at low intensity) due to other factors.....
33 UV SpectroscopyInstrumentation – Sample HandlingIn general, UV spectra are recorded solution-phaseCells can be made of plastic, glass or quartzOnly quartz is transparent in the full nm range; plastic and glass are only suitable for visible spectraConcentration is empirically determinedA typical sample cell (commonly called a cuvet):
34 UV SpectroscopyInstrumentation – Sample HandlingSolvents must be transparent in the region to be observed; the wavelength where a solvent is no longer transparent is referred to as the cutoffSince spectra are only obtained up to 200 nm, solvents typically only need to lack conjugated p systems or carbonylsCommon solvents and cutoffs:acetonitrile 190chloroform 240cyclohexane 1951,4-dioxane 21595% ethanol 205n-hexane 201methanol 205isooctane 195water 190
35 UV Spectroscopy The Spectrum The x-axis of the spectrum is in wavelength; nm for UV, for UV-VIS determinationsDue 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 lmaxlmax = 206 nm252317376AbsWavelength (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 solutionWhere A is absorbancee is the molar absorbtivity with units of L mol-1 cm-1l is the path length of the sample (typically in cm).c is the concentration of the compound in solution,expressed in mol L-1A = elc = log I0/Ie= intensity of the incident lightTransmitted light= intensity of the transmitted lightIncident lightl = width of the cuvetteA = 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, AFrom 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 effectthe greater the concentration (c) of the sample, the more UV light will be absorbed – linear effectsome electronic transitions are more effective at the absorption of photon than others – molar absorptivity, e --this may vary by orders of magnitude…eA = elc = log I0/I
38 UV SpectroscopyThe SpectrumThese effects are combined into the Beer-Lambert Law: A = e c lfor 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 Mmolar absorptivities vary by orders of magnitude:values of are termed high intensity absorptionsvalues of are termed low intensity absorptionsvalues of 0 to 103 are the absorptions of forbidden transitionsA is unitless, so the units for e are cm-1 · M-1 and are rarely expressedSince 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*spns*p*alkanescarbonylsunsaturated cmpds.O, N, S, halogensp*Energynps
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 transitionss* (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 ~ nme ~ ( not strong)n p*, p p* requires unsaturated functional groupsmost commonly used, energy good range for UV/Vis ~ nmn p* : e ~p p* : e ~ 1000 – 10,000
41 UV Spectroscopy: Chromophores DefinitionRemember the electrons present in organic molecules are involved in covalent bonds or lone pairs of electrons on hetero-atoms such as O or NSince 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 nms*s
43 UV Spectroscopy: Organic Chromophores Alcohols, ethers, amines and sulfur compounds – in these compoundsthe 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 LUMOs*CNnN sp3sCN
44 UV Spectroscopy: Chromophores Alcohols, ethers, amines and sulfur compoundsn* transition lower in energy than σ*n* transition - - between 150 and 250 nm.max maxH2OCH3OHCH3ClCH3I(CH3)2S(CH3)2OCH3NH(CH3)3Nn* transitionExplain 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, respectivelyEven though this transition is of lower energy than s s*, it is still in the far UV – however, the transition energy is sensitive to substitutionp* ~ nmp
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 carbonylsThis transition is also sensitive to substituents on the carbonylSimilar 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*npσ σ* transitions omitted for clarity
49 UV Spectroscopy: Organic Chromophores hvnppssThe 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: n max = 188 nm ; = 1860 (intense band) n max = 279 nm ; = 15
50 UV Spectroscopy: Chromophores n* and * TransitionsMost UV/vis spectra involve these transitions.* are generally more intense than n*max max typeC6H13CH=CH *C5H11CC–CH *OCH3CCH n*CH3COH n*CH3NO n*CH3N=NCH n*
51 Absorption Characteristics of Some Common Chromophores ExampleSolventlmax (nm)emaxType of transitionAlkenen-Heptane17713,000pp*AlkyneC5H11CC-CH317819622510,0002,000160_Carbonyln-Hexane1862801802931,00016Large12ns*np*CarboxylEthanol20441AmidoWater21460Azo3395NitroCH3NO2Isooctane22NitrosoC4H9NOEthyl ether30066510020NitrateC2H5ONO2Dioxane270
52 UV Spectroscopy: Chromophores Substituent EffectsThe attachment of substituent groups (other than H) can shift the energy of the transitionAuxoxhromes: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 chromophoreBathochromic shift (red shift) – a shift to longer l; lower energyHypsochromic shift (blue shift) – shift to shorter l; higher energyHyperchromic effect – an increase in intensityHypochromic effect – a decrease in intensityHyperchromiceHypsochromicBathochromicHypochromic200 nm700 nm
54 UV Spectroscopy: Substituent Effects Conjugation – most efficient means of bringing about a bathochromic and hyperchromic shift of an unsaturated chromophore:lmax nm e175 15,000217 21,000258 35,000,000n p*p p*n p*p p* ,100