1. Optical Electronic Spectroscopy 1
Important note on lecture design: X-ray methods are also obviously a form of “electronic” spectroscopy, but typically involve higher energies (inner electrons), and are outside the UV-visible region. As such they require somewhat different instrumentation. This lecture focuses on the types of electronic spectrometry that can be accomplished with optical components and radiation.
Lecture Date: January 23rd, 2008

2. The Electromagnetic Spectrum
The visible region of the spectrum comprises photon energies of 36 to 72 kcal/mole. The near UV region (to 200 nm) extends this energy range to 143 kcal/mole.
UV-Visible
X-ray

3. What is Electronic Spectroscopy?
Spectroscopy of the electrons surrounding an atom or a molecule: electron energy-level transitions
Atoms: electrons are in hydrogen-like orbitals (s, p, d, f)
Molecules: electrons are in molecular orbitals (HOMO, LUMO, …)
From
(The Bohr model for nitrogen)
(The LUMO of benzene)

4. Optical Electronic Spectroscopy
Definition: Spectroscopy in the optical (UV-Visible) range involving electronic energy levels excited by electromagnetic radiation (often valence electrons).
This lecture is related to the “high-energy” (“non-optical”) electron spectroscopy covered in the X-ray lecture
Methods:
Atomic absorption
Atomic emission (e.g ICP-OES)
Molecular UV-Visible absorption
Luminescence, Fluorescence, Phosphorescence
Important note on lecture design: X-ray methods are also obviously a form of “electronic” spectroscopy, but typically involve higher energies (inner electrons), and are outside the UV-visible region. As such they require somewhat different instrumentation. This lecture focuses on the types of electronic spectroscopy that can be accomplished with optical components and radiation.

5. Definitions of Electronic Processes
Emission: radiation produced by excited molecules, ions, or atoms as they relax to lower energy levels.
Absorption: radiation selectively absorbed by molecules, ions, or atoms, accompanied by their excitation (or promotion) to a more energetic state.
Luminescence: radiation produced by a chemical reaction or internal electronic process, possibly following absorption.

6. More Electronic Processes
Fluorescence: absorption of radiation to an excited state, followed by emission of radiation to a lower state of the same multiplicity
Occurs about 10-5 to 10-8 seconds after photon absorption
Phosphorescence: absorption of radiation to an excited state, followed by emission of radiation to a lower state of different multiplicity
Occurs about 10 to 10-5 seconds after photon absorption
Fluorescence is a process that involves a singlet to singlet transition.
Phosphorescence is a process that involves a triplet to singlet transition

7. What is Emission?
Atoms/molecules are driven to excited states (in this case electronic states), which can relax by emission of radiation.
M + heat  M*
Higher energy
E = hn
Lower energy
Other process can be active, such as “non-radiative” relaxation (e.g. transfer of energy by random collisions).
M*  M + heat
OES = Optical Emission Spectroscopy

8. What is Absorption?
Electromagnetic radiation travels fastest in a vacuum.
When EM radiation travels through a substance, it can be slowed by propagation “interactions” that do not cause frequency (energy) changes:
Absorption does involve frequency/energy changes, since the energy of EM radiation is transferred to a substance, usually at specific frequencies corresponding to natural atomic or molecular energies
Absorption occurring at optical frequencies involves low to mid-energy electronic transitions.
c = the speed of light (~3.00 x 108 m/s)
i = the velocity of the radiation in the medium in m/s
ni = the refractive index at the frequency i
For more, see pg. 125 of Skoog et al.

9. Absorption and Transmission
Transmittance:
T = P/P0 P0 P
Absorbance:
A = -log10 T = log10 P0/P b
The Beer-Lambert law commonly uses absorbance as a measure of the absorption rather than %transmittance, because of the linear relationship between concentration and absorbance.
A is linear vs. b!
(A preferred over T)
Graphs from

10. The Beer-Lambert Law The Beer-Lambert Law (a.k.a. Beer’s Law): A = ebc
Where the absorbance A has no units, since A = log10 P0 / P
e is the molar absorbtivity with units of L mol-1 cm-1
b is the path length of the sample in cm
c is the concentration of the compound in solution, expressed in mol L-1 (or M, molarity)
Beer’s law can be derived from a model that considers infinitesimal portions of a “block” absorbing photons in their cross-sections, and integration over the entire block
Beer’s law is derived under the assumption that the fraction of the light absorbed by each thin cross-section of solution is the same
See pp of Skoog, et al. for details
The Beer-Lambert law commonly uses absorbance as a measure of the absorption rather than %transmittance, because of the linear relationship between concentration and absorbance.

11. Deviations From the Beer-Lambert Law
Deviations from Beer’s law (i.e. deviations from the linearity of absorbance vs. concentration):
Intermolecular interactions at higher concentrations
Chemical reactions (species having different spectra)
Peak width/polychromatic radiation
Beer’s law is only strictly valid with single-frequency radiation
Not significant if the bandwidth of the monochromator is less than 1/10 of the half-width of the absorption peak at half-height.
Detailed discussion of Beer’s Law can be found in Skoog et al., Chapter 13.
For an alternative view, see: Bare, William D. A More Pedagogically Sound Treatment of Beer's Law:
A Derivation Based on a Corpuscular-Probability Model, J. Chem. Educ. 2000, 77, 929.

12. Deviations from the Beer-Lambert Law
Intermolecular interactions at higher concentrations:
Dimers,
oligomers
Detailed discussion of Beer’s Law can be found in Skoog et al., Chapter 13.
Figure from Chapter 5 of Cazes, Analytical Instrumentation Handbook 3rd Ed. Marcel-Dekker 2005.

13. Deviations from the Beer-Lambert Law
Deviations caused by use of polychromatic light on a spectrum in which e changes a lot over the bandwidth of the light.
Consider two wavelengths a and b with a and b
 = 1000, 1000
 = 1500, 500
 = 1750, 250
Absorbance (A)
Detailed discussion of Beer’s Law can be found in Skoog et al., Chapter 13.
Concentration (M)

14. Basic Instrument Layout for Optical Spectroscopy
Absorption:
Radiation
Source
Sample
Wavelength
Selector
Detector
(photoelectric transducer)
Fluorescence, Phosphorescence and Scattering:
Sample
Wavelength
Selector
Detector
(photoelectric transducer)
Radiation
source
(90° angle)
See Skoog et al. Figure 7-1.
Emission and chemi-luminescence
Sample
(source)
Wavelength
Selector
Detector
(photoelectric transducer)

15. Atomization: The Dividing Line for Atomic/Molecular
Samples used in optical atomic (elemental) spectroscopy are usually atomized
This destroys molecules (if present) and leaves the atoms
The UV-visible spectrum of the atoms is of interest, not the molecular spectrum.
In the case of X-ray methods – the X-ray spectrum is not sensitive to molecular effects and atomization is not necessary! This is because core electrons are usually affected by the higher-energy X-ray radiation – physical and chemical state have only minute effects on these core electrons (i.e. gas vs solid, oxide vs. element).

16. Elemental Analysis
Elemental analysis – qualitative or quantitative determination of the elemental composition of a sample
Optical electronic methods are heavily used in elemental analysis
Other elemental analysis methods not discussed here:
Mass spectrometry (MS), e.g. ICP-MS
X-ray methods
Other methods (radiochemical)
Classical
Note: there are many other methods for elemental analysis which are not based on optical electronic spectroscopy.

17. Atomic Electronic Energy Levels
Electronic energy level transitions in hydrogen – the simplest of all!
Balmer series (visible)
Transitions start (absorption) or end (emission) with the first excited state of hydrogen
Lyman series (UV)
Transitions start (absorption) or end (emission) with the ground state of hydrogen
In 1885 Balmer derived a simple equation relating the emission lines for hydrogen. It is: 1/ = R(1/22 – 1/n2)
R is the Rydberg constant x 107 m-1
The equation can be generalized for the
Lyman series: 1/ = R(1/22 – 1/n2) Paschen series:
Brackett series:
Pfund series:
Diagrams from

18. Atomic Electronic Energy Levels
Term symbols and electronic states: used to precisely define the state of electrons
spin
multiplicity
s = total spin quantum number
j = total angular momentum quantum number
l = orbital quantum number (s,p,d,f…)
mj = state
Term: 2P
s,p,d,f,g
(l value)
Level:
2P3/2
2j+1
State:
2P3/2-1/2
Used to denote energy levels, and label term (Grotrian) diagrams for the hydrogen atom
Electronic configuration (1s22s1 for Li, etc…) does not give all of the info about the arrangement of electrons in a atom.
Ex. 2p2, where the two electrons could occupy orbitals with different angular momentum (when l=1, ml can be 1, 0, and –1). Also no spin is given by 2p2.
Figure from the Sapphire Electronic Spectroscopy Software Package, Cavendish Instruments Limited.

19. Atomic Electronic Energy Levels
Term symbol (Grotrian) diagram for the sodium atom
Each transition on the diagram can be linked to a peak in the spectrum
The number of lines can approach 5000 for transition-metal elements.
Line broadening can be caused by:
Doppler effects
pressure broadening (collisions)
Lifetime of state (uncertainty)
Figure from H. A. Strobel and W. R Heineman, Chemical Instrumentation: A Systematic Approach, Wiley, 1989.

20. Atomic Electronic Energy Levels
The population of energy levels partly determines the intensity of an emission peak
The Boltzmann distribution relates the energy difference between the levels, temperature, and population:
E = energy of state
P = number of states having equal energy at each level
N = number of atoms in state
Key point: to get more atoms into excited states, you need higher temperatures. (See example 8-2, problem 8-9)
Element/Line (nm)
Ne/Ng at 2000 K
Ne/Ng at 3000 K
Ne/Ng at K
Na 589.0
9.9 x 10-6
5.9 x 10-4
2.6 x 10-1
Ca 422.7
1.2 x 10-7
3.7 x 10-5
1.0 x 10-2
Zn 213.8
7.3 x 10-15
5.4 x 10-10
3.6 x 10-3
(Values from Cazes pg 79, Table 1)

21. Atomic Electronic Energy Levels
The simulated spectrum for the sodium atom

22. Atomic Emission Two types of emission spectra: Examples: Continuum
Line spectra
Examples:
ICP-OES (inductively-coupled plasma optical emission spectroscopy), also known as ICP-AES
LIBS (laser-induced breakdown spectroscopy)
Note that emission here refers to spontaneous emission, not the stimulated emission used by lasers.

23. Torches and Atomic Emission
History: Emission came first (study of sunlight by Fraunhofer in 1817, identification of spectral “lines”), studied throughout the 1800’s and early 1900’s
Before the use of the plasma for OES in 1964, the flame/gas torch (or arc/spark, etc…) had the following problems:
Temperature instability
Not hot enough to excite/decompose all materials
Atomizer/
Emission Source
Temperature (°C)
Flame
Plasma (e.g. ICP)
Electric arc
Electric spark
>10000
Today: The plasma has become the almost universally-preferred method
History: atomic emission placed demands on monochromators
Today: Technology has led to polychromators/detectors with sufficient resolution
DCP torches – not as hot, fewer emission lines, much more maintenance.

24. Plasma Torches
Plasma: a low-density gas containing ions and electrons, controlled by EM forces
DCP torches – not as hot, fewer emission lines, much more maintenance.

25. Plasma Torches
In the inductively-coupled plasma (ICP) torch, the sample will reside for several milliseconds at K.
Other torches – direct current plasma
Microwave induced plasma
DCP torches – not as hot, fewer emission lines, much more maintenance.
An argon ICP torch in action:
Photo by Steve Kvech,

26. More on Plasma Torches Another view of an argon ICP torch:
DCP torches – not as hot, fewer emission lines, much more maintenance.
Diagram from Lagalante, Appl. Spect. Reviews. 34, 191 (1999)

27. Arc and Spark Sources for Atomic Emission
Arc and spark sources – used for qualitative analysis of organic and geological samples
Only semi-quantitative because of source instability
Spark sources achieve higher energies
Several mg of solid sample is packed between electrodes, 1-30 A of current is passed achieving several hundred volts potential.
Applications include metals analysis or cases where solids must be analyzed.

28. Atomic Emission: Mono- and Polychromators
Diffraction gratings are used to select wavelengths (in combination with collimating lens, and slits)
Echelle (ladder) gratings: high dispersion and high resolution
~ grooves/mm typical for UV-Vis work
Require filters to isolate “orders” (i.e. n=1)
m  = d(sin i + sin r)

29. Atomic Emission: Detectors
At the end of the spectrometer, photons are detected.
Commonly used detectors:
Photomultiplier tubes (PMT) – dynamic range 109
Solid-state detectors:
Charge-coupled devices (CCD) – 1D or 2D arrays (charge readout or “transfer” devices)
Silicon photodiodes with thousands of individual elements
Very sensitive, very well-suited to echelle grating polychromators, very fast

30. Modern ICP-OES Spectrometers
Example system:
Varian Vista PRO
Features:
1. Axial flame view
2. Echelle grating polychromator
3. CCD detector
CCD chips are often made of sub-arrays matched to emission lines.
This design features an axial view of the plasma, a echelle polychromator, a CCD chip detector, and no moving parts.
It can measure wavelengths between nm simultaneously. Its advantages translate into the following benefits: less time and sample needed for multiple element scans, the ability to measure multiple wavelengths for a given element for confirmation in quantitative analyses, and reduced flicker noise because the line, background, and/or internal standard can be measured all at the same time.
Typical cost: $110K USD
Figure from Varian Vista PRO sales literature.

31. Detection Limits of ICP-OES
Typical detection limits (Varian Vista MPX):
Considerations include the number of emission lines, spectral overlap
Linearity can span several orders of magnitude.
See also Figure in Skoog, et al.
Note – axial view maximizes sensitivity, radial view maximizes stability.
Note limits in ppb assuming equal densities.

32. Atomic Absorption – Early History
In the beginning – atomic emission was the only way to do elemental analysis via optical spectroscopy
Bunsen and Kirchhoff (1861) – invented a non-luminous flame to study emission. Showed that alkali elements in the flame removed lines from a continuous source.
Walsh (1955) – notices that molecular spectra are often obtained in absorption (e.g. UV-Vis and IR), but atomic spectra are always obtained in emission. Proposes to use atomic absorption (AA or AAS) for elemental analysis
Advantages over emission – far less interference, avoids problems with flame temperature
Electrothermal vaporizers/atomizers also used, take advantage of conductive heating.

33. Atomic Absorption (AA) and Elemental Analysis
Atomic absorption spectrometry is one of the most widely used methods for elemental analysis.
Basic principles of AA:
The sample is atomized via:
A flame (methane/H2/acetylene and air/oxygen)
An electrothermal atomizer (an electrically-heated graphite tube or cup)
UV-Visible light is projected through the flame
The atoms absorb light (electronic excitation), reducing the beam
The difference in intensity is measured by the spectrometer
Source P0
Sample/Flame P
Monochromator
Dual flame/electrothermal atomizers are now available. Electrothermal vaporizers/atomizers are used for higher sensitivity applications, take advantage of conductive heating. See chapter 9 of Skoog et al.
Detector
Images are of Aurora AI1200,

34. Atomic Absorption: Sources
Hollow cathode lamps – sputtering of an element of interest, generating a line emission spectrum:
EDL stands for electrode-less discharge lamp.
Anode – the element in an electron tube, sometimes called the plate, that attracts the electrons emitted by the cathode.
Cathode – the element of an electron tube that emits electrons.
Typical linewidths of nm (0.02Å)
Other AA Sources: electrode-less discharge lamp (EDL) – see Skoog Ch 9B-1

35. Atomic Absorption: Monochromators
The monochromator filters out undesired light in AA (typical bandwidths are 1 angstrom/0.1 nm)
Unlike ICP-OES, where the mono- or polychromator actually analyzes the frequency.
In other words – there is no need to scan the grating, just set (aimed through a slit) and run
Echelle (ladder) gratings are popular:
Figure from T. Wang, in J. Cazes, ed, “Ewing’s Analytical Instrumentation Handbook”

36. Other Features of Atomic Absorption Systems
Sample nebulizers: Produces aerosols of samples to introduce into the flame (oxyacetylene is the hottest)
Detectors: Common examples are photomultiplier tubes, CCD (charge-coupled devices), and many more.
Monochromator: removes emissions from the flame (flame is often kept cool just to avoid emission)
Modulated source (chopper): also removes the remaining emissions from the flame. The signal of interest is given an AC modulation and passed through a high-pass filter.
Spectral interferences:
Absorption from other things (besides the element of interest) – other flame components, particulates, etc… Scattering can cause similar problems
Background correction can help

37. Detection Limits of Atomic Absorption Systems

38. How Are Elements Actually Analyzed?
For AA and ICP-OES, samples are dissolved or digested into solution.
Samples are flowed into the flame/plasma and analyzed.
Two methods for quantitative analysis:
Standard calibration: the unknown sample’s absorbance/emission is compared with several references which “bracket” the expected concentration. (Linear relationship)
Standard addition: the unknown sample is divided into several portions. One portion is directly analyzed, the others have the reference material added in varying amounts. The linear relationship is determined, and the intercept is used to calculate the real concentration of the unknown
At the end: the results yield elements in ppm, ppb, mg/mL, etc…

39. Atomic Fluorescence
Developed as an alternative to AA and ICP-OES, with potentially greater sensitivity.
Has not yet achieved widespread use but cheaper tunable lasers may change this.
Laser – stimulated emission (coherent emission from an excited state induced by a second photon)
Processes: hv
Non-radiative
Thermal hv hv
Non-radiative hv
Resonance
Direct Line
Stepwise
Thermally-assisted

40. (photoelectric transducer)
Atomic Fluorescence
Instrumentation
Sample
Wavelength
Selector
Detector
(photoelectric transducer)
Radiation
source
(90° angle)
Sources include hollow-cathode lamps, electrodeless discharge tubes (brighter), and lasers (brightest)
EDLs produce atomic line spectra that is brighter than HCTs. Ionized argon in a sealed quartz tube is used to excite a small quantity of a metal (or salt). The argon is excited and ionized by strong RF at ~20-40 MHz.
Picture from Perkin-Elmer

41. Laser-Induced Breakdown Spectroscopy (LIBS)
Just like ICP-OES, except a focussed laser creates the plasma:
Fiber optic
Figure from US Army/Ames

42. Elemental Analysis with Optical Spectroscopy
A comparison of the techniques – the choice is not always clear!
Plasma Emission (ICP-OES)
AA (Flame)
Atomic Fluorescence
Dynamic Range
Wide
Limited
Qualitative Analysis
Good
Poor
Multielement Scan?
Trace Analysis
Small samples
Matrix interferences
Low
High
Spectral interferences
Cost
Moderate
Speciated analysis: The analysis of atomic “species”, elements in chemically distinguishable environments.
Examples of hyphenation to add “speciation”:
ICP-OES coupled to a HPLC
AA coupled to a GC

43. Homework Problems Optical Electronic Spectroscopy
Chapter 8:
Problem 8-9
Chapter 10:
Problem 10-2

44. Review Skoog et al. Chapters 6-10 Review Cazes Chapters 3-4
Further Reading
Review Skoog et al. Chapters 6-10
Review Cazes Chapters 3-4
Optical Electronic Spectroscopy
H. A. Strobel and W. R. Heineman, “Chemical Instrumentation: A Systematic Approach”, 3rd Ed., Wiley (1989).