READ: Energy Levels  

In Bohr’s model of an atom, negative electrons circle the positive nucleus at different energy levels –(fixed distances from the nucleus where the electrons are found)-. The model below shows an atom of the element nitrogen using Bohr’s model. Bohr’s model is useful for understanding properties of elements and their chemical interactions, but more modern and accurate models of the atom demonstrate that electrons do not orbit the nucleus in fixed paths, they actually occupy regions of space called orbitals.    


Electrons absorb energy and get pushed up energy levels. When they drop energy levels they emit energy in the form of light. The greater the drop in energy levels, the greater the energy of the photon of light given off by the electron. Because the color of light depends on its energy, we see different colors depending on the energy change of the electron.    

 

The light emitted by the sign containing neon gas (on the left) is different from the light emitted by the sign containing argon gas (on the right).  

Each Element Has a Unique Spectrum   

The light frequencies emitted by atoms are mixed together by our eyes so that we see a blended color. Several physicists, including Anders J. Angstrom in 1868 and Johann J. Balmer in 1875, passed the light from energized atoms through glass prisms in such a way that the light was spread out so they could see the individual frequencies that made up the light.  


In the figure above, we see the emission spectrum, or atomic spectrum, for hydrogen gas. The emission spectrum of a chemical element is the unique pattern of light obtained when the element is subjected to heat or electricity.   When hydrogen gas is placed into a tube and electric current passed through it, the color of emitted light is pink. But when the light is separated into individual colors, we see that the hydrogen spectrum is composed of four individual frequencies. The pink color of the tube is the result of our eyes blending the four colors. Every atom has its own characteristic spectrum; no two atomic spectra are alike. The image below shows the emission spectrum of iron. Because each element has a unique emission spectrum, elements can be identified using them.  



You may have heard or read about scientists discussing what elements are present in the sun or some more distant star and wondered how scientists could know what elements are present in a place no one has ever been. Scientists determine what elements are present in distant stars by analyzing the light that comes from those stars and using the atomic spectrum to identify the elements emitting that light.  In the same way that we can identify elements by their line emission spectrum, elements can also be identified using a flame test.  In a flame test an element is burned in a flame and the results flame color is indicative and unique to the element.    

The results of some flame tests are as follows: 

  • Lithium burns red 
  • Copper burns green/blue 
  • Magnesium burns white 
  • Strontium burns red 
  • Barium burns yellow/green 
  • Potassium burns light purple 
  • Sodium burns yellow/orange.   

Energy Levels  

The key idea in Bohr’s model of the atom is that electrons occupy definite orbits that require the electron to have a specific amount of energy. In order for an electron to be in the electron cloud of an atom, it must be in one of the allowable orbits and it must have the precise energy required for that orbit. Orbits closer to the nucleus would require smaller amounts of energy for an electron and orbits farther from the nucleus would require the electrons to have a greater amount of energy. The possible orbits are known as energy levels. One of the weaknesses of Bohr’s model was that he could not offer a reason why only certain energy levels or orbits were allowed.   

Bohr hypothesized that the only way electrons could gain or lose energy would be to move from one energy level to another, thus gaining or losing precise amounts of energy. The electrons take quantum leaps –(leaps of specific sizes)- as they move through energy levels.  To understand quantum leaps imagine a ladder that has rungs only at certain heights. The only way you can be on that ladder is to be on one of the rungs and the only way you could move up or down would be to move to one of the other rungs. Suppose we had such a ladder with 10 rungs. Other rules for the ladder are that only one person can be on a rung and in normal state, the ladder occupants must be on the lowest rung available. If the ladder had five people on it, they would be on the lowest five rungs. In this situation, no person could move down because all the lower rungs are full. Bohr worked out rules for the maximum number of electrons that could be in each energy level in his model and required that an atom is in its normal state (ground state) had all electrons in the lowest energy levels available. Under these circumstances, no electron could lose energy because no electron could move down to a lower energy level. In this way, Bohr’s model explained why electrons circling the nucleus did not emit energy and spiral into the nucleus.   

Bohr’s Model and Atomic Spectra   


The evidence used to support Bohr’s model came from the atomic spectra. He suggested that an atomic spectrum is created when the electrons in an atom move between energy levels. The electrons typically are at ground state –(the state closest to the nucleus where electrons have the lowest energy possible)-. If the electrons are given energy (through heat, electricity, light, etc.), the electrons in an atom could absorb energy by jumping to an excited state -(a higher energy state farther from the nucleus)-. When the electrons fall back to ground state, the electrons then give off the absorbed energy in the form of a photon -(a light particle)-. The energy emitted by electrons dropping back to lower energy levels would always be in precise amounts of energy because the differences in energy levels are precise. This explains why you see specific lines of light when looking at an atomic spectrum - each line of light matches a specific “step down” that an electron can take in that atom.

             

http://en.wikipedia.org/wiki/File:Bohr_atom_model_English.svg   

CC-BY-NC-SA Utah State Office of Education. Material adapted from ck12.org


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Last modified: Thursday, 9 June 2016, 2:26 PM