Experiment 2: Atomic Emission Spectra                                                      Chem 111 – Spring ‘05

 

Introduction

            Atoms consist of electrons surrounding a very dense nucleus made of protons and in most cases neutrons. The number of protons in the nucleus is what distinguishes different types of atoms (i.e. elements), and isotopes are atoms whose nuclei consist of the same number of protons, but different numbers of neutrons.  In neutral atoms, the number of electrons surrounding the nucleus is the same as the number of protons in the nucleus.

            Centuries ago scientists began using flame tests (literally putting compounds into a flame) to identify certain elements in those compounds based on the color of the flame that resulted. As it turns out, when the light from a colored flame is observed through a prism, rather than a continuous rainbow of colors, sharp lines of individual colors appear. These lines are called emission spectra of atoms.  This was an unexpected observation and it puzzled scientists for many decades. 

 In the early years of the 20th century Niels Bohr developed a theory of the structure of atoms that included the critical idea that the electrons in atoms can only have specific energy values, which he called energy levels.  If electrons are ‘excited’ meaning they have acquired additional energy, they can only be excited to the allowed energy levels within that specific atom. Furthermore, if these excited electrons released their acquired energy as photons of light, they could only release light of certain energies associated with the specific energy levels.  This theory helped explain the observed lines of color in the emission spectra of excited atoms.  The lines of the emission spectrum of a certain atom are associated with allowed electronic transitions between energy levels in the atom. The figure below is a diagram of the energy levels of a hydrogen atom, and the arrows pointing downward represent allowed electronic transitions from higher energy levels to lower energy levels. Three of the sets of arrows are given the names Lyman, Balmer, and Pachen. These sets of arrows, or series of lines, are named after the scientist who discovered them.  As it turns out, only one of these sets of emission lines is in the visible region of the electromagnetic spectrum.

Experimental Procedure

Part 1

 

Hydrogen emission spectrum: Using the fiber optic and spectrometer – laptop set-up you used during week 1, collect the Hydrogen emission spectrum and print out a copy of the spectrum for you and your lab partner.  Make sure you collect a spectrum that does not include any peaks of ‘stray’ or background light (the room will be dark when you’re collecting the spectra).

  1. On your printed spectrum, note the colors and intensities of the peaks and put this spectrum into your lab notebook.
  2. Using a diffraction grating (designed to spread out the light), see if you can confirm – with your eyes - the number of lines in your spectrum in the visible region. 
  3. Using the equation below where λ = wavelength in m, nf = the final (or lower) energy level the electron, and ni = the initial (or higher) energy level of the electron, complete following calculations and enter them in a similar table in your lab notebook.

An example calculation is provided:

            nf =1; ni =2

           

nf

ni

λ (in m)

λ (in nm)

In the visible spectrum? yes or no

If yes, what color is this light?

1

2

1.215x10-7nm

121.5 nm

no

-

1

3

 

 

 

 

1

4

 

 

 

 

2

3

 

 

 

 

2

4

 

 

 

 

2

5

 

 

 

 

2

6

 

 

 

 

3

4

 

 

 

 

3

5

 

 

 

 

 

Part 1 Questions: Answer these in your lab notebook

  1. According to figure 1 (previous page), which series, Lyman, Balmer or Paschen, is the series of lines in the visible spectrum? 
  2. Could you ‘see’ all of these lines? Were they all the same intensity?

 

Additional emission spectra: He, Hg, Na

 

Carefully observe the other elemental emission spectra in the laboratory using both the fiber optic and the diffraction grating.

  1. For two of the additional samples, collect and print the emission spectrum from the laptop and note the ID number of the ‘tube’ of gas for each ‘unknown’ you choose.  Label the spectrum with the colors of the most intense lines based on the wavelengths and based on your own observations with the diffraction grating. Put the two spectra into your notebook.
  2. Using available text books or on-line resources (http://onsager.bd.psu.edu/~jircitano/periodic4.html) assign the spectra you observed to either He, Hg, or Na.  Make sure you provide a proper citation for the sources you use.

 

Conclusions

Write some general conclusions in your lab notebook based on your observations and what you learned from this experiment.

 

For additional background information on atomic emission spectra, see:

http://csep10.phys.utk.edu/astr162/lect/light/absorption.html