Bohr Model Lab

Name: ______Date:______Block: ______Bohr Model Lab

Introduction: In 1912 scientists didn’t understand much about the electron structure of the atom, but they knew this was the key to understanding reactivity since it was clear that the gain and loss of electrons was involved in reactions. At this point in history, Rutherford had done his famous gold foil experiment and discovered the nucleus, but no one understood why, if the positive charge was located in the center, the negatively charge electrons didn’t fall into the nucleus.

Around this time, scientists were also discovering the quantum nature of energy. Phenomena like blackbody radiation and the photoelectric effect led scientists to believe that energy was transferred between bodies in certain sized packets called quanta. Niels Bohr used this idea of a quantum nature of energy coupled with the unique atomic emission spectra to conclude that electrons must exist in specific energy levels. In other words, electrons were not floating around the electron cloud in random paths, but were orbiting in discrete orbital paths around the nucleus. His model was easy to understand because it is similar to how the planets orbit the sun and Bohr is credited with what people call the planetary model of the atom. Bohr won the Nobel Prize for his work, and it’s important to understand that his ideas were not just guesses, but were grounded in mathematics that provided a way to make theoretical predictions that could be verified by experimental evidence. He developed an equation based on theory that ultimately was able to quantitatively predict the wavelengths given off in hydrogen’s spectra and even hydrogen’s ionization energy. Although his model was limited to Hydrogen, the development of a mathematical model able to make real-world predictions is what made his ideas so influential.

Bohr’s model is based on the idea that electrons inhabit discrete energy levels. When an electron gains energy, it moves to an excited state (another energy level farther away from the nucleus). Eventually, the electron loses that energy and moves back to its ground state. When it loses that energy, we see it emitted as visible light. Because the energy levels are discrete, electrons can only absorb and emit certain amounts of energy. Bohr put together the ideas of Max Planck who found that the energy of a given frequency of radiation could be calculated by E=hv, where E = energy in Joules, h = Planck’s constant (6.626x10-34 J-s) and v = frequency in Herz and the ideas of J.J. Balmer who found a mathematical relationship between the wavelengths of the spectral lines of a hydrogen atom.

Purpose: In this lab you are going to observe Hydrogen’s spectral lines and then calculate the energy of each of these transitions. From this you will be able to use Bohr’s equation (see number 3 below) to calculate which energy level the electron is traveling from to give off each spectral line. (You will be assessed on your Data Collection & Processing - DCP) Procedures: 1. Observe the hydrogen emission lines through a spectroscope 2. Record the colors and wavelengths of the lines (one significant figure is not sufficient).

Analysis: You should show your work, organize it into tables on a separate page and attach to this page. 1. Convert the wavelengths of light to frequencies (c= v where c is the speed of light 3.00 x 108 m/s). Then calculate the energy of each color of light. (E = hv, where h is Planck’s constant 6.626 x 10-34 J-s). You can do this in two separate steps or one step if you combine the equations. Name: ______Date:______Block: ______2. Find the energy level transition for each color a. Since the emission lines are in the visible region, these transitions are known as the Balmer series. Balmer series transitions occur when an electron moves from some excited state to the 2nd energy level. These transitions are named after J.J. Balmer because he was the first to describe the lines according to the following equation: 1  1 1   R     22 n 2 

Where n =3,4,5… and R = Rydberg constant of 1.1 x 107 m-1. Balmer did not know at the time that this equation represented the electrons transitioning from one energy level to another. It was Bohr who combined the work of Rydberg and Planck to come up with his model and changed Balmer’s equation slightly to:

骣 1 1 DE = R '琪 - 琪 2 2 桫nlower n upper

where nlower is the ground state energy level and nupper is the excited state energy level and R = -2.178x10-18 J. 3. Since you probably did not get a whole number in number 3, compare to the nearest whole number that makes sense to calculate your percent error for each transition. **As always: put your final results into a clearly labeled table 4. Calculate the ionization energy of the hydrogen atom. (Hint: see pg 301 sample exercise 7.5) 5. Explain your calculation from #5 6. The Paschen series occur when electrons transition from some excited state to the 3rd energy level. Calculate one of the wavelengths of light given off in these transitions and determine the region in the electromagnetic spectrum (e.g. microwave, infrared, visible, etc.) 7. The Lyman series occur when electrons transition from some excited state to the 1st energy level. Calculate one of the wavelengths of light given off in these transitions and determine the region in the electromagnetic spectrum (e.g. microwave, infrared, visible, etc.) 8. Is the energy given off in a transition more or less for a transition between energy level 5 to 4 or from 2 to 1? Explain. Use this to predict and draw a diagram of how the spacing between energy levels changes as one gets farther from the nucleus. 9. Explain why emission spectra can be used to identify unknown elements. 10. Suggest a composition for the unknown below. (There are at least three elements that are not part of the substance.) Name: ______Date:______Block: ______

11. How could scientists use this method to determine the composition of stars? 12. Distinguish between emission, absorption and continuous spectra.