Radio-waves
Region
Frequency
(HZ)
10 6 - 10 10
Wavelength 10m – 1 cm
Microwaves
Region
10 10 - 10 12
Infra-red
Region
10 12 - 10 14
1 cm – 100µm 100µm – 1µm
Energy
NMR, ESR
0.001
J/mole
– 10
Rotational
Spectroscopy
Vibrational spectroscopy
Order of some
100 J/mole
Some 10 4
J/mole
Visible
Region
10 14 - 10 15
Ultra-violet
Region
10 15 - 10 16
700 – 400 nm 400-10 nm
X-ray
Region
10 16 - 10 18

-ray
Region
10 18 - 10 20
10nm
100pm
– 100pm
1 pm
–
Electronic
Spectroscopy
Electronic
Spec.
Some 100 kJ/mole
Some 100s kJ/mole
10 7 - 10
J/mole
9 10 9 - 10
J/mole
11
Frequency (

)
Wavelength (

)

Energy of light
Frequency of light

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



-rays X-rays UV
Visible
IR Microwave Radio
Longer Wavelength, Lower Energy

 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

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.

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.

or
Continuous Energy Loss Quantized Energy Loss

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

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

2
Line Emission Spectrum of Hydrogen Atoms

 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.

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.

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

 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* p*
Unoccupied levels
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: i) s
, p
, n electrons
Sigma ( s
) – single bond electron
A MO A

y y

UV Spectroscopy
Pi ( p
) – double bond electron
High energy anti-bonding orbital ( π*)
Low energy bonding orbital ( π)
Nonbonding 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 p n n p * s * p * s * p * alkanes carbonyls unsaturated cmpds.
O, N, S, halogens carbonyls

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: p n n s s p * s * p * s * p * alkanes carbonyls unsaturated cmpds.
O, N, S, halogens carbonyls
150 nm
170 nm
180 nm √ - if conjugated!
190 nm
300 nm √
30

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 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
NH
2
O O
 max
= 206 nm
252
317
376
35
Wavelength (nm)

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 =
Incident light e
lc = log I
0
/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
Transmitted light
= intensity of the 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 = 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 others – path length absorbed – linear effect
 the greater the
( l ) through the sample will cause more UV light to be concentration absorbed – linear effect
 some electronic transitions are more effective at the absorption of photon than molar absorptivity, e
( c

) of the sample, the more UV light will be this may vary by orders of magnitude…
A = e
lc = log I
0
/I e

UV Spectroscopy
The Spectrum
These effects are combined into the Beer-Lambert Law: A = e c l
 for most UV spectrometers,
1 cm in path length) l would remain constant (standard cells are typically
 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
• values of 10 3 -10 4 are termed high intensity absorptions low intensity absorptions
• values of 0 to 10 3 are the absorptions of
A is unitless, so the units for e forbidden transitions 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 expressed as e directly or as the logarithm of e.
e , and the y-axis is

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:
Energy s* p* n p s s s p n n p * s * p * s * p * alkanes carbonyls unsaturated cmpds.
O, N, S, halogens 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 n (non-bonding) ss
* pp
* n
s
* n
p
* p
(bonding) s
(bonding) ss
* transition in vacuum UV (

~ 150 nm) n
s
* 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 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: Organic Chromophores
 Alkanes
150 nm
(CH
4
, C
2
H
6 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),  ~
C C s* s
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 like the alkane s  s * transition also possible.
 value ( < 200 nm );
Note how this transition occurs from the HOMO to the LUMO s*
CN
C N n
N sp 3
C N C N anitbonding
orbital s
CN
C N

UV Spectroscopy: Chromophores
Alcohols, ethers, amines and sulfur compounds n
s
* transition lower in energy than σ s
* n
s
* transition -  between 150 and 250 nm.
H
2
CH
CH
CH
O
(CH
3
3
3
(CH
3
(CH
3
3
OH
Cl
I
)
)
)
2
2
3
S
O
CH
3
NH
N
2
 max
167
184
173
258
229
184
215
227
Explain why  max and the corresponding e max is different.
e max
1480
150
200
365
140
2520
600
900

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

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

C O
UV Spectroscopy: Organic Chromophores n  p *
Carbonyls – unsaturated systems incorporating N or O can undergo transitions ( ~285 nm ) in addition to p  p * n  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

C O
UV Spectroscopy: Organic Chromophores p* s* p* s* n p* n 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 s*  max n p*  max
= 188 nm ; e = 1860
(intense band)
= 279 nm ; e = 15

UV Spectroscopy: Chromophores n p * and pp * Transitions
Most UV/vis spectra involve these transitions. pp * are generally more intense than n p *
C
6
H
13
CH=CH
2
C
5
H
11
C  C–CH
3
O
CH
3
CCH
3
O
CH
3
COH
CH
3
NO
2
CH
3
N=NCH
3
 max
177
178
186
204
280
339 e max type
13000 pp *
10000 pp *
1000 n s *
41
22
5 n p * n p * n p *

Absorption Characteristics of Some Common Chromophores
Chromophore Example Solvent
 max
(nm) e max
Alkene
Alkyne
Carbonyl
Carboxyl
Amido
Azo
Nitro
Nitroso
Nitrate
C
6
H
13
HC CH
2
C
5
H
11
C

C-CH
3
n-Heptane
n-Heptane
O
CH
3
CCH
3
O
CH
3
CH
O
CH
3
COH
O
CH
3
CNH
2
H
3
CN NCH
3
CH
3
NO
2
C
4
H
9
NO
C
2
H
5
ONO
2
n-Hexane
n-Hexane
Ethanol
Water
Ethanol
Isooctane
Ethyl ether
Dioxane
339
280
300
665
270
177
178
196
225
186
280
180
293
204
214
13,000
10,000
2,000
160
1,000
16
Large
12
41
60
5
22
100
20
12
Type of transition p  p
* p  p
*
_
_ n  s
* n  p
* n  s
* n  p
* n  p
* n  p
* n  p
* n  p
*
_ n  p
* n  p
*

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

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

UV Spectroscopy: Substituent Effects
Conjugation – most efficient means of bringing about a bathochromic and hyperchromic shift of an unsaturated chromophore:
H
2
C
CH
2
 max
175 nm e
15,000
217
258
21,000
35,000
465 125,000

-carotene
O
O n  p * 280 p  p * 189 n  p * 280 p  p * 213
12
900
27
7,100