A Absorption spectrum It is that product of the interaction of light with a material or substance in any of its physical states. But the definition goes beyond a simple visible light, since the interaction includes a wide segment of the range of wavelengths and energy of electromagnetic radiation..
Therefore, some solids, liquids or gases, can absorb photons of different energies or wavelengths; from ultraviolet radiation, followed by visible light, to radiation or infrared light, stepping into microwave wavelengths.
The human eye only perceives the interactions of matter with visible light. Likewise, it is able to contemplate the diffraction of white light through a prism or a medium in its colorful components (upper image).
If the ray of light were "caught" after having traveled through a material, and analyzed, the absence of certain bands of colors would be found; that is, black stripes would be observed contrasting with its background. This is the absorption spectrum, and its analysis is fundamental in instrumental analytical chemistry and astronomy..
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The upper image shows a typical absorption spectrum of elements or atoms. Note that the black bands represent the absorbed wavelengths, while the others are the emitted ones. This means that, in contrast, an atomic emission spectrum would look like a black band with stripes of emitted colors..
But what are these stripes? How to know in short if the atoms absorb or emit (without introducing fluorescence or phosphorescence)? The answers lie in the allowed electronic states of the atoms.
Electrons are able to move away from the nucleus leaving it positively charged while they transit from a lower energy orbital to a higher energy one. For this, explained by quantum physics, they absorb photons of a specific energy to carry out said electronic transition..
Therefore, the energy is quantized, and they will not absorb half or three quarters of a photon, but rather specific frequency values (ν) or wavelengths (λ).
Once the electron is excited, it does not remain for unlimited time in the electronic state of higher energy; releases the energy in the form of a photon, and the atom returns to its basal or original state.
Depending on whether the absorbed photons are recorded, an absorption spectrum will be obtained; and if the emitted photons are recorded, then the result will be an emission spectrum.
This phenomenon can be observed experimentally if gaseous or atomized samples of an element are heated. In astronomy, by comparing these spectra, the composition of a star can be known, and even its location relative to the Earth..
As can be seen in the first two images, the visible spectrum includes colors from violet to red and all their shades regarding how much the material absorbs (dark nuances).
The wavelengths of red light correspond to values from 650 nm onwards (until they disappear in infrared radiation). And on the extreme left, the violet and purple tones cover the wavelength values up to 450 nm. The visible spectrum then ranges from 400 to 700 nm approximately.
As λ increases, the frequency of the photon decreases, and therefore its energy. Thus, violet light has higher energy (shorter wavelengths) than red light (longer wavelengths). Therefore, a material that absorbs purple light implies electronic transitions of higher energies..
And if the material absorbs the color violet, what color will it reflect? It will appear a greenish-yellow color, which means that its electrons make very energetic transitions; while if the material absorbs the lower energy red color, it will reflect a bluish-green color.
When an atom is very stable, it generally exhibits very distant electronic states in energy; and therefore you will need to absorb higher energy photons to allow electronic transitions:
Molecules have atoms, and these also absorb electromagnetic radiation; however, their electrons are part of the chemical bond, so their transitions are different. One of the great triumphs of the molecular orbital theory is the power to relate the absorption spectra with the chemical structure.
Thus, single, double, triple, conjugated bonds, and aromatic structures, have their own electronic states; and therefore absorb very specific photons.
By having several atoms, in addition to intermolecular interactions, and the vibrations of their bonds (which also absorb energy), the absorption spectra of the molecules have the form of "mountains", which indicate the bands that comprise the wavelengths where electronic transitions occur.
Thanks to these spectra, a compound can be characterized, identified, and even, through multivariate analysis, quantified.
The upper image shows the spectrum of the methylene blue indicator. As its name evidently indicates, it is blue; but, can it be verified with its absorption spectrum?
Note that there are bands between the wavelengths of 200 and 300 nm. Between 400 and 500 nm there is almost no absorption, that is, it does not absorb violet, blue, or green colors.
However, it has a strong absorption band after 600 nm, and therefore has low-energy electronic transitions that absorb photons of red light..
Consequently, and given the high values of molar absorptivities, methylene blue exhibits an intense blue color..
As can be seen in the image, the green line corresponds to the absorption spectrum of chlorophyll a, while the blue line corresponds to that of chlorophyll b.
First, the bands where the molar absorptivities are greatest must be compared; in this case, those to the left, between 400 and 500 nm. Chlorophyll a absorbs purple colors intensely, while chlorophyll b (blue line) does so with blue color.
By absorbing chlorophyll b around 460 nm, the blue, the yellow color is reflected. On the other hand, it also absorbs strongly near 650 nm, orange light, which means that it exhibits the color blue. If yellow and blue mix, what is the result? The color green.
And finally, chlorophyll a absorbs the blue-violet color, and also a red light near 660 nm. Therefore, it exhibits a green color “softened” by yellow..
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