1. Absorption of EM radiation

2. Molecular absorption processes
~10-18 J
Electronic transitions
UV and visible wavelengths
Molecular vibrations
Thermal infrared wavelengths
Molecular rotations
Microwave and far-IR wavelengths
Each of these processes is quantized
Translational kinetic energy of molecules is unquantized
Increasing energy
~10-23 J

3. Atomic and molecular vibrations correspond to excited energy levels in quantum mechanics
Energy levels are everything in quantum mechanics.
Excited level
DE = hn
Energy
Ground level
The atom is vibrating at frequency, ν.
The atom is at least partially in an excited state.
For a given frequency of radiation, there is only one value of quantum energy for the photons of that radiation
Transitions between energy levels occur by absorption, emission and stimulated emission of photons

4. Excited atoms emit photons spontaneously
When an atom in an excited state falls to a lower energy level, it emits a photon of light.
Excited level
Energy
Ground level
Molecules typically remain excited for no longer than a few nanoseconds. This is often also called fluorescence or, when it takes longer, phosphorescence.

5. Absorption spectra of molecules
Hypothetical molecule with three allowed energy levels
Note relationship to emission!
νij = ΔEij/h
allowed transitions
positions of the absorption lines in the spectrum of the molecule
Line positions are determined by the energy changes of allowed transitions
Line strengths are determined by the fraction of molecules that are in a particular initial state required for a transition
Multiple degenerate transitions with the same energy may combine

6. Fluorescence
Fluorescent lighting exploits this phenomenon: certain phosphors emit visible light when bombarded with UV light. Much more efficient than incandescent lighting.
Also whitening agents in detergents...

7. Atoms and molecules can also absorb photons, making a transition from a lower level to a more excited one
Excited level
This photon has been absorbed
Energy
Ground level

8. In 1916, Einstein showed that another process, stimulated emission, can occur
Before After
Spontaneous emission
Absorption
Stimulated emission

9. Stimulated emission

10. Interaction of radiation with matter
Wavelength
If there are no available quantized energy levels matching the quantum energy of the incident radiation, then the material will be transparent to that radiation

11. X-ray interactions
Quantum energies of x-ray photons are too high to be absorbed by electronic transitions in most atoms - only possible result is complete removal of an electron from an atom
Hence all x-rays are ionizing radiation
If all the x-ray energy is given to an electron, it is called photoionization
If part of the energy is given to an electron and the remainder to a lower energy photon, it is called Compton scattering

12. Ultraviolet interactions
Near UV radiation (just shorter than visible wavelengths) is absorbed very strongly in the surface layer of the skin by electron transitions
At higher energies, ionization energies for many molecules are reached and the more dangerous photoionization processes occur
Sunburn is primarily an effect of UV radiation, and ionization produces the risk of skin cancer

13. UV SO2 and O3 absorption spectra
Absorption cross-section represents the probability that a photon will be absorbed by a molecule of gas. TOMS Wavelengths not perfectly placed for SO2 as it is an Ozone instrument. Wavelengths are UV, Huggins bands. Spectral resolution affects our ability to resolve SO2 band structure and hence SO2 sensitivity and noise.

14. Ultraviolet interactions
UV-A: nm
UV-B: nm
UV-C: nm
The ozone layer in the upper atmosphere absorbs most harmful UV radiation before it reaches the surface
Higher UV-B frequencies are ionizing radiation and can produce harmful effects such as skin cancer
The ionosphere is a region of the upper atmosphere ionized by solar radiation

15. Visible light interactions
Visible light is also absorbed by electron transitions
Higher energies at blue wavelengths relative to red wavelengths: hence red light is less strongly absorbed than blue light
Absorption of visible light causes heating, but not ionization
Car windshields transmit visible light but absorb higher UV frequencies

16. Infrared (IR) interactions
Quantum energy of IR photons ( eV) matches the ranges of energies separating quantum states of molecular vibrations
Vibrations arise as molecular bonds are not rigid but behave like springs

17. Microwave interactions
Quantum energy of microwave photons ( eV) matches the ranges of energies separating quantum states of molecular rotations and torsion
Note that rotational motion of molecules is quantized, like electronic and vibrational transitions  associated absorption/emission lines
Absorption of microwave radiation causes heating due to increased molecular rotational activity
Most matter transparent to µ-waves, microwave ovens use high intensity µ-waves to heat material

18. Molecular dipole moments
For a molecule to absorb IR radiation it must undergo a net change in dipole moment as a result of vibrational or rotational motion.
The electric dipole moment for a pair of opposite charges of magnitude q is the magnitude of the charge times the distance between them, with direction towards the positive charge.
The total charge on a molecule is zero, but the nature of chemical bonds is such that positive and negative charges do not completely overlap in most molecules. Such molecules are said to be polar because they possess a permanent electric dipole moment.
Water is a good example of a polar molecule:
Molecules with mirror symmetry like oxygen, nitrogen and carbon dioxide have no permanent dipole moments.

19. Molecular polarizability
The polarizability of an atom or a molecule is a measure of the ease with which the electrons and nuclei can be displaced from their average positions (e.g., by an external electric field)
When the electrons occupy a large volume of space, e.g., in an atom or molecule with many electrons, the polarizability of the substance is large. When an atom or molecule has large polarizability the magnitude of the instantaneous dipole can be large.
The polarizability of molecules is important in Raman spectroscopy, based on Raman scattering.

20. Key atmospheric constituents
Diatomic, homonuclear molecules (e.g., N2, O2) have no permanent electric dipole moment (also CO2)
Molecular N2, the most abundant atmospheric constituent, has no rotational absorption spectrum
Oxygen (O2) has rotational absorption bands at 60 and 118 GHz
Linear and spherical top molecules have the fewest distinct modes of rotation, and hence the simplest absorption spectra
Asymmetric top molecules have the richest set of possible transitions, and the most complex spectra
Note lack of permanent electric dipole moment in CO2 and CH4 No

21. Vibration modes of simple molecules
Fundamental or normal modes
Symmetric stretch
Bend (Scissoring)
Asymmetric stretch
A normal mode is IR-active if the dipole moment changes during mode motion.
Overtones, combinations and differences of fundamental vibrations are also possible (e.g., 2v1, v1+v3 etc.)
A non-linear molecule of N atoms has 3N-6 normal modes of vibration; a linear molecule has 3N-5.

22. Absorption frequency for a diatomic molecule
m1, m2 = atomic mass of vibrating atoms
c = speed of light [3×108 m s-1]
V = wavenumber [cm-1]
Av = Avogadro’s number [6.023×1023 atoms mole-1]
k = force constant (bond strength) [dynes cm-1]
For a single bond, k = 5×105 dynes cm-1
For a double bond, k = 10×105 dynes cm-1
For a triple bond, k = 15×105 dynes cm-1

23. Infrared (IR) interactions
Vibrational transitions are associated with larger energies than ‘pure’ rotational transitions.
Vibrations can be subdivided into two classes, depending on whether the bond length or angle is changing:
Stretching (symmetric and asymmetric)
Bending (scissoring, rocking, wagging and twisting)
Stretching frequencies are higher than corresponding bending frequencies (it is easier to bend a bond than to stretch or compress it)

Bonds to hydrogen have higher stretching frequencies than those to heavier atoms.

Triple bonds have higher stretching frequencies than corresponding double bonds, which in turn have higher frequencies than single bonds

24. Infrared (IR) interactions
Region
Wavelength [µm]
Energy [meV]
Wavenumber [cm-1]
Type of excitation
Far IR
10 – 200
Lattice vibrations,
Molecular rotations
Mid IR
Molecular vibrations
Near IR
Overtones

25. Absorption spectra of molecules
V = Vibrational quantum number
J = Rotational quantum number
Electronic, vibrational and rotational energy levels are superimposed
The absorption spectrum of a molecule is determined by all allowed transitions between pairs of energy levels, and whether the molecule exhibits a sufficiently strong electric or magnetic dipole moment (permanent or otherwise) to interact with the radiation field

26. Vibrational-rotational transitions
P branch (ΔJ = -1)
Q branch (ΔJ = 0)
(pure vibration)
R branch (ΔJ = +1)
Relative positions of transitions in the absorption spectrum of a molecule

27. Rotational absorption spectrum
Photon frequency associated with a rotational transition

28. Hydrogen chloride (HCl) spectrum
Q branch (ΔJ = 0)
P branch
R branch
Vibrational-rotational absorption spectrum of HCl: shows affect of two chlorine isotopes with slightly different mass

29. Transmittance spectrum for ozone (O3)

30. Transmittance spectrum for CO2

31. Transmittance spectrum for H2O

32. Absorption line shapes
Doppler broadening: random translational motions of individual molecules in any gas leads to Doppler shift of absorption and emission wavelengths (important in upper atmosphere)
Pressure broadening: collisions between molecules randomly disrupt natural transitions between energy states, so that absorption and emission occur at wavelengths that deviate from the natural line position (important in troposphere and lower stratosphere)
Line broadening closes gaps between closely spaced absorption lines, so that the atmosphere becomes opaque over a continuous wavelength range.

33. Pressure broadening
Absorption coefficient of O2 in the microwave band near 60 GHz at two different pressures. Pressure broadening at 1000 mb obliterates the absorption line structure.

34. Rovibrational Energy
Vibrational and rotational transitions usually occur simultaneously splitting up vibrational absorption lines into a family of closely spaced lines
Rotational energy also dependent on direction of oscillation of dipole moment relative to axis of symmetry
When oscillates parallel, ΔJ = 0 transition is forbidden and only P and R branches are seen
When oscillates perpendicular, P, Q and R branches are all seen
The rotational constant is not the same in different vibrational states due to a slight change in bond-length, and so rotational lines are not evenly spaced in a vibrational band
Rovibrational transitions in a CO2 molecule
Diagram taken from Patel (1968)

35. Sulfur dioxide (SO2) ν1: 1151 cm-1, 8.6 µm ν3: 1361 cm-1, 7.3 µm

36. Sulfur dioxide (SO2)
ν1+ν3: 2500 cm-1, 4 µm

37. Water vapor (H2O) Most important IR absorber
Asymmetric top → Nonlinear, triatomic molecule has complex line structure, no simple pattern
3 vibrational fundamental modes
Higher order vibrational transitions (Δv >1) give weak absorption bands at shorter wavelengths in the shortwave bands
2H isotope (0.03% in atmosphere) and 18O (0.2%) adds new (weak) lines to vibrational spectrum
3 rotational modes (J1, J2, J3)
Overtones and combinations of rotational and vibrational transitions lead to several more weak absorption bands in the NIR o o H H
symmetric stretch
v1 = 2.74 μm
bend
v2 = 6.25 μm
asymmetric stretch
v3 = 2.66 μm

38. Transmission spectrum of H2O

39. Explain the peaks in ni….

40. Absorption Spectrum of H2O
v1=2.74 μm
v2=6.25 μm
v3=2.66 μm

41. Carbon dioxide (CO2)
Linear → no permanent dipole moment, no pure rotational spectrum
Fundamental modes:
The v3 vibration is a parallel band (dipole moment oscillates parallel to symmetric axis), transition ΔJ = 0 is forbidden, no Q branch, greater total intensity than v2 fundamental
The v2 vibration is perpendicular band, has P, Q, and R branch
The v3 fundamental is the strongest vibrational band, but the v2 fundamental is most effective due to “matching” of vibrational frequencies with terrestrial Planck emission function
13C isotope (1% of C in atmosphere) and 17/18O isotope (0.2%) cause a weak splitting of rotational and vibrational lines in the CO2 spectrum o c o
symmetric stretch
v1 = 7.5 μm => IR inactive
asymmetric stretch
v3 = 4.3 μm
bend
v2 = 15 μm
bend v2

42. IR Absorption Spectrum of CO2v3 v2

43. Which is the most potent greenhouse gas?
Top: SF6 (sulfur hexafluoride); global warming potential ~23000 times that of CO2 over a 100 year period. SF6 is extremely long-lived as it is inert (lifetime of years)

44. Ozone (O3)
Ozone is primarily present in the stratosphere except anthropogenic ozone pollution which exists in the troposphere
Asymmetric top → similar absorption spectrum to H2O due to similar configuration (nonlinear, triatomic)
Strong rotational spectrum of random spaced lines
Fundamental vibrational modes
14.3 μm band masked by CO2 15 μm band
Strong v3 band and moderately strong v1 band are close in frequency, often seen as one band at 9.6 μm
9.6 μm band sits in middle of 8-12 μm H2O window and near peak of terrestrial Planck function
Strong 4.7 μm band but near edge of Planck functions o o o o
symmetric stretch
v1 = 9.01 μm
bend
v2 = 14.3 μm
asymmetric stretch
v3 = 9.6 μm

45. IR Absorption Spectrum of O3
v1/v3 v2

46. Methane (CH4) Spherical top
5 atoms, 3(5) – 6 = 9 fundamental modes of vibration
Due to symmetry of molecule, 5 modes are degenerate, only v3 and v4 fundamentals are IR active
No permanent dipole moment => No pure rotational spectrum
Fundamental modes H C C C C H H H
v4 = 7.7 µm v1 v2
v3 = 3.3 µm

47. IR Absorption Spectrum of CH4v3 v4
7.6 µm band in otherwise largely transparent part of atmosphere
Methane concentrations also directly/indirectly affected by human activities

48. Nitrous oxide (N2O)
Linear, asymmetric molecule (has permanent dipole moment)
Has rotational spectrum and 3 fundamentals
Absorption band at 7.8 μm broadens and strengthens methane’s 7.6 μm band.
4.5 μm band less significant as it is at the edge of the Planck function.
Fundamental modes: O N N
symmetric stretch
v1 = 7.8 μm
asymmetric stretch
v3 = 4.5 μm
bend
v2 = 17.0 μm
bend v2

49. IR Absorption Spectrum of N2O
v3=4.5 µm
v1=7.8 µm
v2=17 µm

51. Mineral and rock reflectance spectra
Electronic transitions in solids; Fe2+ (iron) particularly important in remote sensing – minerals contain Fe2+ ions
Fundamental vibrational modes of H2O: 2.74 µm, 6.25µm, 2.66 µm
In rock spectra, whenever water is present we see 2 absorption bands in near-IR spectra – one near 1.45 µm (2ν3 overtone) and one near 1.9 µm (v2+v3 combination). Sharpness of bands relates to sites in crystal structure occupied by the water molecules.
Note that penetration depth into natural surfaces is usually restricted to the upper few microns. Consequences?

52. Geological mapping/prospecting
Escondida Mine, Atacama Desert, Chile
ASTER visible
ASTER short-wave IR (SWIR)

53. Why are most plants green and then red or yellow in the fall?
Chlorophyll absorbs in the red and blue, and hence reflects in the green.
Its absorption spectrum is due to electronic transitions
During spring and summer, leaves get their green cast from chlorophyll, the pigment that plays a major role in capturing sunlight. But the leaves also contain other pigments whose colors are masked during the growing season. In autumn, trees break down their chlorophyll and draw some of the components back into their tissues. Conventional wisdom regards autumn colors as the product of the remaining pigments, which were finally unmasked. In other words, autumn leaves were a tree's gray hair.
But in recent years, scientists have recognized that autumn colors probably play an important role in the life of many trees. Yellow leaves get their color from a class of pigments called carotenoids. Another group of molecules, anthocyanins, produce oranges and reds. Trees need energy to make carotenoids and anthocyanins, but they cannot reclaim that energy because the pigments stay in a leaf when it dies. If the pigments did not help the tree survive, they would be a waste. What's more, leaves actually start producing a lot of new anthocyanin when autumn arrives.
"The reds are not unmasked-they are made in autumn," said Dr. David Lee, a botanist at Florida International University.
Evolutionary biologists and plant physiologists offer two different explanations for why natural selection has made autumn colors so widespread, despite their cost. Dr. William Hamilton, an evolutionary biologist at Oxford University, proposed that bright autumn leaves contain a message: they warn insects to leave them alone.
Dr. Hamilton's "leaf signal" hypothesis grew out of earlier work he had done on the extravagant plumage of birds. He proposed it served as an advertisement from males to females, indicating they had desirable genes. As females evolved a preference for those displays, males evolved more extravagant feathers as they competed for mates.
In the case of trees, Dr. Hamilton proposed that the visual message was sent to insects. In the fall, aphids and other insects choose trees where they will lay their eggs. When the eggs hatch the next spring, the larvae feed on the tree, often with devastating results. A tree can ward off these pests with poisons.
Photo and notes text borrowed from the NY Times.
Plot borrowed from Peter v. Sengbusch -
Comment from Ron Fox regarding the missing green absorption of chlorophyll:
First off for the visual system, rod cells have an absorption max at 510 nm, i.e in the green. The three cone cells have maxima at 445, 545 and 585. All peaks are broad but the one at 545 is called the green cone. Thus the visual system has evolved to use what the plants don't use. Blue-green algae and red algae have pigments that do use the green ( ) missed by chlorophyll. The solar spectrum today at sea level is monotonically decreasing from a high at 700 to a low at 400, as far as the visible is concerned. The green part is less than the red part but by only a bit. A younger sun that was cooler would still be strong in the red but less so in the green. However, there is debate about the young sun's surface temperature. Chlorophyll a and chlorophyll b have complementary absorptions in separate portions of the red, suggesting that evolution tried to do as well as possible with the red.
In the fall, trees produce carotenoids, which reflect yellow, and anthocyanins, which reflect orange and red.

54. Why Mars looks red
Iron oxides prevalent in Martian soil show increased reflectance at the red end of the visible spectrum.

55. WTC dust spectra Chrysotile standard
Spectra of World Trade Center dust
Gypsum (water-bearing) = pulverized wallboard
Chrysotile = asbestos (fireproof coatings?)
OH- (hydroxyl ion) – has first overtone at ~1.4 µm – most common feature present in near-IR spectra of terrestrial materials

56. Spectral response of vegetation
Reflectance of vegetation in Visible – SWIR region

57. Light excites atoms, which emit light that adds (or subtracts) with the input light.
When light of frequency w excites an atom with resonant frequency w0:
Electric field at atom
Electron cloud
Emitted field + =
Incident light
Emitted light
Transmitted light
On resonance (w = w0)
An excited atom vibrates at the frequency of the light that excited it and re-emits the energy as light of that frequency.
The crucial issue is the relative phase of the incident light and this re-emitted light. For example, if these two waves are ~180° out of phase, the beam will be attenuated. We call this absorption.

58. Refractive Index vs. Wavelength
Since resonance frequencies exist in many spectral ranges, the refractive index varies in a complex manner.
Electronic resonances usually occur in the UV; vibrational and
rotational resonances occur in the IR; and inner-shell electronic
resonances occur in the x-ray region.
n increases with frequency, except in anomalous dispersion regions.