Atomic Line Spectra
Introduction
In the late 19th century, physicists noticed something about heated elements that they could not explain with their current model of the atom; the Rutherford model of the atom stated only that the electrons travel around the dense nucleus like planets around the sun but did not describe their motion at all. When an element was heated up until it glowed and then that light was directed through a prism the light was only certain wavelengths and not the entire spectrum. Scientists used two different ways to view this spectrum for each element. The emission spectrum of an element is made by heating up the element until it glows and then sending the light through a prism. The absorption spectrum is made by sending light through the gaseous state of the element and then refracting the remaining light through a prism. The emission and absorption spectra appear as shown in figure 1. The emission and absorption spectra are direct inverses of each other. In other words if the two spectra were combined you would have the continuous spectrum.
Figure 1: The difference between the continuous, emission, and absorption spectra.
Scientists also found that each element had an emission spectrum unique to that element. The uniqueness of the spectrum allows scientists to identify elements just by looking at the emission or absorption spectrum. This uniqueness combined with the fact that only certain wavelengths were coming out of the atoms confused the physicists of the late 19th century because the current model of the atom did not explain it. In order to explain this phenomenon Niels Bohr came up with a new model of the atom.
New Explanation
Bohr’s new model of the atom ended up giving the answer to why each element only gave off certain wavelengths of light. Bohr’s new model assumed that electrons could only travel around the nucleus of the atom in certain distinct energy levels. This meant that in order for an electron to orbit the nucleus it had to have one of a certain set of angular momentums. Because of this assumption, Bohr’s model perfectly explained the emission spectrum for hydrogen but did not work as well for atoms of more than one electron because it did not take into account electron-electron repulsions. As an electron moved between these certain defined energy levels around the atom, it would give off a certain wavelength of light. Also it would take a certain amount of energy in order to excite the electron to a new energy level. In figure 2 it shows that an electron must absorb energy such as a photon to go from a lower orbit to a higher one and must emit energy to drop to a lower orbit.
Figure 2: Bohr’s model of the atom and electrons moving energy levels.
Bohr’s model turned the physics world on its head because classical physics said that the electron would slowly lose energy and spiral into the nucleus. Bohr’s model led to the invention of the field of quantum mechanics.
Failure of Classical Mechanics:
Classical mechanics was unable to completely explain the behavior of very small particles such as molecules, atoms, subatomic particles, etc. Under classical mechanics, a particle has a defined mass, position, and momentum at any given point in time. Classical (Newtonian) mechanics only describes typical everyday objects such as cars, projectiles, etc. Furthermore, classical applications treat particles and light separately, which is understandable based on the observations before the discovery of atomic particles. This leads to significant errors when using classical physics to answer quantum problems. However, today’s science sheds a brighter light on the subject of quantum particles; it is now known that the laws of classical physics do not apply to smaller particles that move at speeds near the speed of light. With the development of research in quantum principles, atomic bodies are treated as both particles and waves.
To explain to the scientific world why Atomic Absorption Spectroscopy results in discrete lines of color at differing frequencies, we may look at Max Planck’s hypothesis on oscillating charged particles. He postulated that an oscillator only accepts a certain values of energy; that is, putting just any arbitrary value (i.e. any real number) of energy to an oscillator will probably fail to yield a reaction. These values are small compared to classical values, which is why the quantized energy levels appear continuous on a classical scale.
The appearance of atomic line spectra exhibits that energy has to be quantized to yield results. If arbitrary values could fulfill the energy required to excite particles, a spectrograph of any exited particle would show as a continuous band of the visible light spectrum. This discovery proves that classical mechanics are not suitable to fit the particles in the quantum realm. However, since particles only emit light at certain wavelengths, the energy required to excite the particle must be a quantized value; likewise, it emits quantized energy when it falls to a lower, more stable energy level, creating a unique frequency that emits light of a certain color.
Observations of atomic absorption spectroscopy clearly attest to Max Planck’s hypothesis about the quantization of energy. Furthermore, he knew that his hypothesis required him to digress from Newtonian mechanics in order to create a more feasible explanation for this peculiar behavior.
The Error of the Bohr Model:
Niels Bohr’s planetary model of the atom, though important to the development of quantum mechanics, was imperfect. Bohr was under the assumption that the orbital of the electron revolved around the nucleus in a circular, planetary motion, similar to a planet’s orbit around the sun. However, according to Newtonian mechanics, an accelerated electron must emit energy; and if it were indeed losing energy, the electron would theoretically spiral into the nucleus due to electromagnetic forces, considering the attraction of the nucleus’s positive charge and the electron’s negative charge—but this, of course, is not the case. We know that there is ground state due to the fact that energy can be used to excite the electron, raising it from its lowest energy state to a higher energy state.
Another important example of these unexplained observations is that, because Bohr designed his model primarily using the simplest atom, hydrogen, his model could not predict the energy levels of multi-electron atoms. Because atomic line spectra show the frequencies in which light is emitted, we are able to distinguish the change in energy levels for atoms with multiple electrons, especially those that emit radiation in the visible light spectrum. For example, when an excited atom releases energy, the light it emits has a certain frequency and wavelength; using these two pieces of data, scientists are able to find the initial energy level and the final energy level with the use of the equation:
What scientists found confusing was the uniqueness of the spectrum allowed for the identification and the fact that only certain wavelengths were coming out of the atoms.
The new explanation shows results in the aforementioned equation, that the transition between states can be quantified to find the initial and final energy states.
Future Research
Researchers can use atomic line spectra to investigate the radiative transitions and the energy levels of new and existing elements. Knowing transition and energy level data allows a better understanding of interactions at the atomic level. The Atomic Spectra Database lists the atomic line spectra and associated energy levels of most elements. Currently, the database covers the transitions of 99 elements and the energy levels of 56 elements. Data for energy levels in the lanthanide earth series covers only the first five spectra. Clearly, there are more elements whose properties are unreported, presenting a direction for future research.
Check out the interactive periodic table displaying the known atomic line and emission spectra for individual elements at http://jersey.uoregon.edu/vlab/elements/Elements.html
Atomic Spectra Database (National Institute of Standards and Technology-NIST)
http://www.nist.gov/pml/data/asd.cfm
Resources:
1) http://jersey.uoregon.edu/vlab/elements/Elements.html
2) http://www.nist.gov/pml/data/asd.cfm
3) Atomic absoption and Emission Spectra, Atomic Absoprtion and Emission Spectra, September 25, 2011 <http://csep10.phys.utk.edu/astr162/lect/light/absorption.html>.
4) "Atomic Emission Spectra - the Origin of Spectral Lines." Avogadro Web Site. Web. 08 Oct. 2011. <http://www.avogadro.co.uk/light/bohr/spectra.htm>.