Transcript UV Spectroscopy

Electronic Spectroscopy of
molecules
Regions of Electromagnetic Spectrum
Radio-waves
Region
Microwaves
Region
Infra-red
Region
Visible
Region
Ultra-violet
Region
X-ray
Region
-ray
Region
Frequency
(HZ)
106 - 1010
1010 - 1012
1012 - 1014
1014 - 1015
1015 - 1016
1016 - 1018
1018- 1020
Wavelength
10m – 1 cm
1 cm – 100µm
100µm – 1µm
700 – 400 nm
400-10 nm
10nm –
100pm
100pm –
1 pm
NMR, ESR
Rotational
Spectroscopy
Vibrational
spectroscopy
Electronic
Spectroscopy
Electronic
Spec.
0.001 – 10
J/mole
Order of some
100 J/mole
Some 104
J/mole
Some 100
kJ/mole
Some 100s
kJ/mole
107- 109
J/mole
109- 1011
J/mole
Energy
Frequency ()
Wavelength ()
Electromagnetic Radiation
Energy of light
Frequency of light
E = h
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
UV Spectroscopy
-rays
X-rays
UV
IR
Microwave
Radio
Visible
Longer Wavelength, Lower Energy
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
UltraViolet range
Visible range
150 nm 200 nm
Longer Wavelength, Lower Energy
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.
Flame Test
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.
Emission of Energy
(2 Possibilities)
or
Continuous Energy Loss
Quantized Energy Loss
Emission of Energy
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
Emission Spectrum
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
H2 Emission Spectrum
Line Emission Spectrum of Hydrogen Atoms
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.
Regions of Electromagnetic Spectrum
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.
Origin of Electronic Spectra
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
Origin of Electronic Spectra
UV Spectrum of Isoprene
Types of Electrons in Molecules
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)
Observed electronic transitions: graphical representation
s*
Unoccupied levels
p*
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:
Sigma (s) – single bond electron
i) s, p, n electrons
A
MO
A
MO’s Derived From the 2p Orbitals
y
y
UV Spectroscopy
Pi (p) – 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
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*
alkanes
s
p*
carbonyls
p
p*
unsaturated cmpds.
n
s*
O, N, S, halogens
n
p*
carbonyls
Possible Electronic Transitions
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:
s
s*
alkanes
150 nm
s
p*
carbonyls
170 nm
p
p*
unsaturated cmpds.
180 nm
n
s*
O, N, S, halogens
190 nm
n
p*
carbonyls
300 nm
√ - if conjugated!
√
30
UV Spectrum of Isoprene
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 UVVIS 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
NH2
max =
Abs
O
Wavelength (nm)
O
206 nm
252
317
376
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........
A = elc = log I0/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
= intensity of the incident light
Transmitted light
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 = 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
UV Spectroscopy
The Spectrum
These effects are combined into the Beer-Lambert Law:
A=ecl
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.
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*
p*
Energy
n
p
s
s
s*
alkanes
s
p*
carbonyls
p
p*
unsaturated cmpds.
n
s*
O, N, S, halogens
n
p*
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
ss*
pp*
n (non-bonding)
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 NO2 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 (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
C
s*
s
C
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  value (< 200 nm);
like the alkane s  s* transition also possible.
Note how this transition occurs from the HOMO to the LUMO
s*CN
C
N
nN sp
C
3
C
N
N
anitbonding
orbital
sCN
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.
max
emax
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
Explain why max and the corresponding emax is different.
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
UV Spectroscopy: Organic Chromophores
Alkenes
C C
s*
p*
s*
p*
= hv
=hc/
hv
p
p
s
s
p
p*
Example: ethylene absorbs at max = 165 nm e= 10,000 (intense band)
UV Spectroscopy: Organic Chromophores
C O
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
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
UV Spectroscopy: Organic Chromophores
s*
p*
C O
n
p*
n
s*
p*
hv
p
n
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 = 188 nm ; e= 1860 (intense band)
np* max = 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*
max
emax
C6H13CH=CH2
177
13000
pp*
C5H11CC–CH3
178
10000
pp*
186
1000
ns*
CH3COH
204
41
np*
CH3NO2
280
22
np*
CH3N=NCH3
339
5
np*
type
O
CH3CCH3
O
Absorption Characteristics of Some Common Chromophores
Chromophore
Example
Alkene
C6H13HC
Solvent
CH2
Alkyne
O
emax
Type of
transition
n-Heptane
177
13,000
pp*
n-Heptane
178
196
225
10,000
2,000
160
pp*
_
_
n-Hexane
186
280
1,000
16
ns*
np*
n-Hexane
180
293
Large
12
Ethanol
204
41
np*
Water
214
60
np*
Ethanol
339
5
np*
C5H11CC-CH3
Carbonyl
max (nm)
CH3CCH3
O
CH3CH
Carboxyl
Amido
O
CH3COH
O
ns*
np*
CH3CNH2
Azo
H3CN
NCH3
Nitro
CH3NO2
Isooctane
280
22
np*
Nitroso
C4H9NO
Ethyl ether
300
665
100
20
_
np*
270
12
np*
Nitrate
C2H5ONO2
Dioxane
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.)
UV Spectroscopy: Substituent Effects
In General – Substituents may have any of four effects on a chromophore
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
e
Hypsochromic
Bathochromic
Hypochromic
200 nm
Hyperchromic
i.
700 nm
UV Spectroscopy: Substituent Effects
Conjugation – most efficient means of bringing about a bathochromic and
hyperchromic shift of an unsaturated chromophore:
H2C
max nm
CH2
e
175
15,000
217
21,000
258
35,000
465
125,000
-carotene
O
n  p* 280
12
900
n  p* 280
27
7,100
p  p* 189
O
p  p* 213