; Absorption Spectroscopy
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Absorption Spectroscopy


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8.1 Absorption Spectroscopy (UV and visible)

1. Introduction
2. Theory
3. Methods
4. Transitions and chromophores
5. Applications
6. CD

1. Introduction
One of the main goals of biochemical (and related) research is to understand the
essence of biological life on a molecular level. Many experimental (physical) methods
provide valuable information such as chromatography or centrifugation. The most
detailed data, including structure information at atomic level, stems however from
spectroscopic methods, i.e. from the interaction of electromagnetic radiation with
biological molecules.
Three types of spectroscopy can be distinguished: absorption (respectively
transmission), emission and scattering. One of the most important methods of the
latter type is x-ray crystallography. A typical example of emission spectroscopy is
fluorescence. Here, we discuss absorption spectroscopy with a focus on radiation
from the visible and UV spectrum. The principle is simple: radiation of a given
frequency interacts with the sample (usually in solution) and the transmitted radiation
is studied.
The following figure summarizes frequency, wavelength and energy ranges of
absorption spectroscopy.

Electromagnetic spectrum with wavelengths from 1 nm to 1 m

    x-rays    UV vis.       IR                micro                 radio
nucl. electron              molecular         electron              nuclear
trns. transitions           vibrations        spins                 spins
Möss-    absorption         IR/Raman          EPR                   NMR
bauer   x-ray UV/vis.

        1      2     3     4      5     6      7            8        9
                     wavelength (nm; powers of 10)

        17    16    15      14    13      12    11     10       9        8
                         frequency (s-1; powers of 10)

        7      6     5      4      3     2      1           0        -1
                     energy (J/mol; powers of 10)

2. Theory
Electromagnetic radiation is characterized by its frequency  and its wavelength ,
which are connected by
       = c                       with c = 2.998*108 m/s.
Visible light with a wavelength of 300 nm has a frequency of 1015 s-1. The energy of a
photon is
      E=h                          where h = 6.63*10-34 Js.
A photon with = 300 nm, has an energy of 6.63*10-19 J.
Various energy units are used in spectroscopy:
- With 1 eV = 1.602*10-19 J the above energy corresponds to 4.139 eV.
- In terms of "wave numbers" this energy is (for  in nm) 107/ = 3.3*104 cm-1.
- Finally, one may indicate the energy of one mol of photons:
       6.02*1023*6.63*10-19 = 3.99*105 J/mol.
For comparison, consider the following energies at room temperature T=300ºK:
- kT: 0.026 eV ~ 4.1*10-21 J ~ 160 cm-1 (per molecule)         k = 1.381 10-23 J/ºK
- RT: 2.5 kJ/mol                                               R = 8.314 J/(ºK mol)
When energy from electromagnetic radiation is transferred to a molecule, the latter is
brought to an excited state. Thus, the energy difference E between the two states of
the molecule must correspond to the incoming radiation energy:
      E = h .
Electronic transitions in a molecule occur with a certain probability. These depend on
selection rules, which depend on symmetry considerations, spin conservation etc. Due
to alterations in the system, induced for example by internal vibrations, symmetries
are not perfect and as a consequence "forbidden" transitions become possible
(although remain weak). Absorption also depends on the populations of the different
states, which in turn assume in an equilibrium situation a Boltzmann distribution, i.e.
they are proportional to exp(E/kT). In an absorption spectrum, peaks indicate the
energy values for which radiation interacts with the sample. The line widths of these
peaks provide an energy interval that arises, besides from instrument artifacts and
alike, from Heisenberg's uncertainty rule: E  h/ with E the energy uncertainty and
 the lifetime of a state.
Consider light with a certain frequency and intensity I0 that enters a sample, i.e. a
solution containing the molecule under investigation at a given concentration C (in
mol/m3 or mM). The probability of absorption by the molecules is proportional to the
number of molecules encountered. Passing a layer of the solution of thickness dx, the
intensity change dI becomes
      dI = - const I0 C dx                   const is a proportionality constant.
Integration over the thickness of the sample d yields I = I0 exp(const C d);
introducing the absorbance or extinction coefficient  = const / ln(10) one writes:
      I = I0 10  C d             Lambert-Beer's law.
In the case of d=1 cm the exponent C d is called optical density OD. It is of
advantage to choose the concentration such that OD is about 0.5.

3. Methods
A typical spectrometer consists of the following elements:
-   a light source, e.g. a conventional lamp for visible light or a deuterium lamp for
-   a monochromator to obtain the desired frequency (or to sweep through an interval
    of frequencies);
-   a rotating mirror (or similar) to split the beam into two parts;
-   two cuvettes with solvent+solute and solvent only, each in the path of one beam;
-   detectors for each beam;
-   an electronic subtraction device for the two beam intensities and a plotter (or


    source        monochromator         mirrors     sample +     detectors   subtraction
                                                    reference                plotting

4. Chromophores
In the context of biological molecules, chromophores (i.e. atomic groups that
efficiently absorb electromagnetic radiation) are often related to ring systems.
Examples are porphyrins in hemoglobins, cytochromes or in chlorophyll. High
extinction coefficients are also observed for coenzymes such as NADH or FADH2.
Another example is the retinal in rhodopsins. On a more general ground one can
observe the aromatic amino acids Phe, Tyr and in particular Trp as well as the bases
of nucleotides. Finally, absorption measurements of the peptide bonds between amino
acids in proteins are sensitive tools of the local conformation. The following table
summarizes some chromophores, their wavelength of maximal absorption and the
extinction coefficient  at this wavelength.

Chromophor              max        at max        comment
                        [nm]        [mol-1 cm-1]
Peptide   *         ~190            ~7000
Peptide n  *         ~220             ~100         "forbidden"
Phe                     257              200
                        206             9300
                        188            60000
Tyr                     274             1400
                        222             8000
                        193            48000
Trp                     280             5600
                        219            47000
Adenosine               259            14900
Guanosine               276             9000
Cytidine                271             9100
Thymidine               267             9700
Uridine                 261            10100
FAD                     438            60000
Chlorophyll a           435            60000         from plants
                        675            85000

Chlorophyll: a simple estimation
Porphyrin groups occur in many biological systems, examples are hemoglobin
(oxygen transport) and chlorophyll (photosynthesis). The following figures show the
chemical structure of chlorophyll a and its absorption spectrum.

In an approximate calculation of chlorophyll we first note that porphyrines have 11
double bounds and thus 22 -electrons. Two of these bonds can be broken without
major change of the absorption spectrum. Thus, 18 -electrons form a ring of de-
localized electrons. Assuming a "circular potential box" with radius r where r is
an average radius of the ring (~4Å), we can write a Hamiltonian for these electrons:
            H =  h2/(2m)·2 =  h2/(2m)·1/r2·/2
With the boundary condition r = |n| we get the solutions
            () = 1/(2)½·exp(in)             n = …-2,-1,0,1,2…
with energies
            En = n2h2/(2mr2).
The 18 electrons fill orbitals with n=0,±1,±2,±3,±4 (two electrons in each). Thus the
lowest energy transition is
            E = h2/(2mr2) * (52-42)
yielding a wavelength of 571 nm, which lies between the observed two absorption
bands (see table and figure above). The introduction of a periodic potential with the
nuclei as wells and of electron-electron interactions splits the spectral line by 0.9 eV,
which corresponds to a splitting of about 240 nm, and thus approximates the
chlorophyll spectrum quite well.

5. Applications
The following are three examples where absorption spectroscopy was applied to
biological systems.

(A) Poly-lysine
Poly-lysine adopts different conformations under different conditions. At high pH it
forms regular secondary structures: an -helix at room temperature and a -sheet at
high temperature. At neural pH the molecules have no defined conformation (random
coil). The following figure shows the extinction coefficient of poly-lysine at pH 10.8
and 25º C (thick line), at pH 10.8 and 52 ºC (dashed line) and at pH 6 and 25 ºC (thin
line). Thus, absorption spectroscopy may serve to characterize the secondary structure
of a molecule.

(B) Environment-dependent absorption
The following figure shows the change in extinction of Trp when transferring the bee
venom melittin (27 amino acids) from water to lipid membranes. It proves the tight
interaction of melittin with membranes, and thus its toxicity.

(C) Extinction of adenosine as a function of ordering
When different forms of adenosine are subjected to absorption spectroscopy one
observes a dependence with the degree of ordering. The following figure (left)
displays the spectra for AMP (adenosine-mono-phosphate), single-stranded polyA
with limited ordering, and a double helix with high ordering. The increase of
absorption with increasing ordering is called hyperchromism. It allows observation of
thermal denaturing of DNA, e.g. with different contents of GC (see figure on the

6. A short introduction to CD
Light is considered (in classic theory) to be a wave with an electric and a magnetic
field component, E and H, respectively. In linear polarized light the E-vector
oscillates in one direction perpendicular to the propagation direction (and so is the H-
vector). As the figure below shows, one can describe a linear polarized wave E as the
sum of two circular polarized waves EL and ER, with the field vectors rotating around
the propagation axis. (Similarly, circular polarized light can be described as a
superposition of two linear polarized waves with different phases.)

      EL    ER

Interaction with molecules can be different for the two circular polarized waves. For
example, chiral molecules can absorb the two waves to different extents. The result is
then a wave with |EL|  |ER|, i.e. an elliptic wave. In Circular Dichroism (CD)
experiments one measures the difference in extinction coefficients L - R for the
two circular polarized waves. One reports the ellipticity  = k··c·d with sample
thickness d (cm), concentration c (g/cm-3) and  in cm2/g; k = ln(10)·180/(2·)=33.
Note the somewhat strange units due to historical reasons.
In CD, optically active molecules must fulfill two conditions. A chromophore is
needed to get absorption and a chiral center is required for  0. The two centers
may be identical, overlap or be located very close two each other. Chirality may be
due to an asymmetric atom such as tetrahedral carbon atoms with four different
substituents. Chirality may also be caused by asymmetry in the structure, e.g. the
spiral-like structure of composite rings or helical structures in proteins or DNA.
Important in all cases is an asymmetry of the electron distribution.
CD is a sensitive and easy-to-use tool to follow conformational changes in proteins
and DNA, e.g. as a consequence of temperature or pH changes. The following figure
shows its use to quantitatively determine the extent of regular secondary structure in
the (unknown) 3D structure of a protein.

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