The following are tips on how I use these images in my classes to introduce the topic of electronic structure of atoms. If you do not have access to a fiber optic spectrometer to perform the demonstration yourself, you can click on the thumbnail images to download a high resolution version to used in your classes.
I start the lesson by introducing spectra with an incandescent light bulb so that students see how the images are correlated to each other. Then I show the spectrum of a hydrogen bulb like that below. The wavelengths of the emission lines are reported in many General Chemistry textbooks, but in this demonstration, they can be directly read off the graph. The peaks appear at 656 nm, 486 nm, 434 nm, and 410 nm.
I use the narrow, well-separated lines in this spectrum to introduce Bohr's model of the atom and the idea of transitions between fixed energy levels leading to emission or absorbance of light at a few well-defined wavelengths. The 656 nm line corresponds to the n = 3 to n = 2 transition, and I ask the students to calculate the wavelength of this transition so that we can confirm that it appears in the spectrum. The following references give background information about the Bohr model and suggest give some hints for teaching this concept at the introductory level.
After they are comfortable with the Bohr model of the hydrogen atom, I show my students the spectrum of helium, like that shown below. They will immediately note that it looks different than the spectrum of hydrogen. I ask them why they think this is so. Hopefully someone will suggest that helium has one more proton than hydrogen, which will attract the electrons more strongly. I show them how to modify Bohr's energy equation to include the atomic number and we recalculate the energy of the n = 3 to n = 2 transition. The resulting wavelength isn't found in the spectrum. After some more thought, one of them will usually suggest that the helium atom also has two electrons. At this point I discuss the fact that we can't know where both electrons are at any particular time (an introduction to the Heisenberg Uncertainty Principle), and that we will need a more complicated theory than Bohr's to explain the appearance of the helium spectrum. This is my starting point for talking about quantum mechanics and orbitals.
The following reference describes an appealing demonstration to show why the positions of the electrons can't be determined at any given time, which prevents us from calculating energies by classical methods:
The reference below describes a laboratory exercise which can be used in an advanced course to correlate the observed emission spectrum of an element with the effective nuclear charge when discussing a more advanced model of atomic structure.
These Supplemental Materials accompany the article "Simultaneous Display of Spectral Images and Graphs Using a Web Camera and Fiber-Optic Spectrophotometer" Niece, Brian K. J. Chem. Educ., 2006, 83, 761-764 and are provided with the permission of the Journal.