Introduction

Atomic line specra are experimentally observed phenomena resulting from the emission or absorption of a photon from the excited state of electrons transitioning between energy levels. There are two types of atomic line spectra, emission and absorption lines. Emission, absorption and Einstein coefficients are used in equations related to atomic line spectra.

Questions 2,3

More importantly, why were atomic line spectra a demonstration of the failures of classical physics?

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.

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.

Why was Bohr’s Model incorrect and how was it corrected in light of what we know about

atomic line spectra?

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:

Future Research 

Researchers can use atomic line spectra to investigate new elements or chemical bonds of interest.

Further Resources

http://jersey.uoregon.edu/vlab/elements/Elements.html

http://www.nist.gov/pml/data/asd.cfm