State-of-the-Art and Trends in Atomic Absorption Spectrometry [chapter]

Hlcio Jos Izrio Filho, Rodrigo Fernando dos Santos Salazar, Maria da Rosa Capri, ngelo Capri, Marco Aurlio Kondracki de Alcntara, Andr Lus de Castro Peixoto
2012 Atomic Absorption Spectroscopy  
Atomic Absorption Spectroscopy 14 concluded that the Fraunhofer lines were produced due to the presence of atomic vapors in the solar atmosphere that absorbed part of radiation emitted by Sun. From 1859 to 1861, Robert Bunsen and Gustov Kirchoff demonstrated that each chemical element had a characteristic color or spectrum when heated to incandescence (Na yellow; K violet). Heating several elements during a flame test, they identified characteristic spectra of these elements and established a
more » ... lation between emission spectrum and absorption spectrum. This explains the black lines in the solar spectrum: atoms in the solar corona absorbing part of energy emitted by Sun (continuous spectrum), originating the black lines observed. This also permits to identify absorbing atoms normally present in the corona, comparing the black lines to elements' emission spectrum produced in the laboratory. At the beginning of the 20th century, the development of quantum mechanics theory provided mathematical patterns explaining the phenomena, which is the interaction of radiation with the matter. Considering this point, the necessary theoretical basis to develop a new technique of elementary analysis using the phenomenon of atomic absorption was established, but only in 1955, the first proposal of a practical instrument was introduced. At the beginning of the 60s, the first commercial equipment appeared and so far, equipments have been using the same basic components, although with more technologically involved. Understanding this technique requires understanding the phenomenon of emission and absorption of radiant-energy through the matter or, more specifically, through atoms. Basic principles Wave-particle duality: Light Electromagnetic radiation is a form of energy described by classical physics as a wave made up of mutually perpendicular, fluctuating electric and magnetic field that propagates at a constant speed. It is characterized by wavelength ( ), (distance between two adjacent waves) and by frequency v (number of waves per unit of time). The speed of light propagated in a vacuum is C=299.792.458 m s -1 . This model explains the energy propagation, but cannot explain its interaction with the matter. That interaction can be explained if we treat energy as a particle, called photon. A photon is characterized by frequency and wavelength and can transfer energy amount E= hv/ , where h is Planck constant (4.135667516×10 −15 eV s). Fig. 1. Regions of the electromagnetic spectrum. Electromagnetic spectrum is conveniently divided in wavelength rates with specific names (from gamma rays to radio waves) as shown in figure 1. That wavelength division does not www.intechopen.com State-of-the-Art and Trends in Atomic Absorption Spectrometry 15 have a physical meaning itself, being only a practical classification in accordance with the available technological equipment for its generation and detection. For example, the visible region of light between 400 and 700 nanometers (nm) is directly detected by the human eye and perceived as visible light (1 nm = 10 -9 m). The atom, energy of a quantum state and electronic transitions Bohr model has an atom consisting of a nucleus containing protons and neutrons surrounded by a cloud of electrons in fact inhabit specific regions in space. This is known as an orbital. The further an orbital gets from the nucleus, the more they gain potential energy associated to a determined orbital. Quantum Mechanics explains that orbitals have quantized energy levels and for moving an electron to another level, it has to receive or emit the exact amount of energy corresponding to the difference between the two electronic levels (ΔE = E 1 -E 0 ). The amount of energy required to move and electron from energy level E 0 to energy level E 1 can be provided by heat due to a collision with other particles or absorb the energy of a photon. In this case, the energy of a Photon (E = h = hc/ ) should be equal to the difference between the orbitals (ΔE), this is, only a Photon of a particular wavelength is absorbed and can promote that transition. This phenomenon is known as atomic absorption. A more stable electron configuration of an atom is the one with less energy, also known as ground state configuration. The difference of energy between the last full orbital and the next empty orbital of the atom in a ground state is of the same order of magnitude of photons with wavelengths between 200 and 800 nm, this means, photons in ultraviolet regions and visible light of electromagnetic spectrum. Sodium atom in ground state, e.g., has an electronic configuration of 1s 2 2s 2 2p 6 3s 1 . The 3s electron can receive a photon with energy of 589.0 nm (E=2.2 eV) and passes to 3p orbital, which is an unstable state known as excited state. Being unstable, the excited atom loses its energy quickly (approximately in 10 -8 s) and returns to ground state. One way to lose excitation energy is by emitting a photon of 589.0 nm, a phenomenon known as atomic emission. A photon of 330.3 nm can also be absorbed by sodium. This is the difference of energy between 3s and a 4p orbital (3.6eV), but one photon of 400 nm cannot be absorbed because there are not two orbitals in a sodium atom with the same difference of energy. The return of electron from 4p orbital to ground state can also occur in two steps: first to 4s and then to 3s by emitting two photons with energies correspondent to the two transitions in a phenomenon known as atomic fluorescence. The sodium atom can also receive enough energy to remove an electron, turning into sodium ion (Na + ), known as ionization. In this case, a change occurs in orbitals of different energy levels, so that the ion has a new set of transition being able to absorb or emit photons of wavelengths differently from metallic sodium. Each chemical element has a unique electronic structure that differentiates from others. This implies in a possible and unique set of transitions, a set of characteristic absorption/emission lines that can be used for identification. Although the set of transitions is unique for each element, there may be a coincidence of spectrum in some lines of two or more elements, which means that different atoms can absorb or emit photons of same wavelength. Even though the theoretical basis was established in the beginning of the twentieth century, only in the early 1950s an Australian physicist, Sir Alan Walsh, proposed the phenomenon
doi:10.5772/26076 fatcat:b2i4szewpfeyxnm5vsarxzgwjq