Investigation 1: Spectroscopy

Heat an element in a flame and its electrons jump up to higher energy levels. They don't stay there. As each electron falls back down, it releases the energy difference as a photon of one exact wavelength. Because every element has its own ladder of energy levels, with its own specific gaps between rungs, the photons it releases land at its own specific wavelengths. That's why sodium burns yellow and strontium burns red: different ladders, different gaps, different colors.

A spectroscope separates those mixed-together photons by wavelength, usually with a diffraction grating or prism, and projects them onto a screen or detector as a series of thin bright lines rather than a smooth rainbow. Each line is one specific transition. The sodium line at 589 nm, for instance, corresponds to an electron falling from the 3p level back down to the 3s ground state in a sodium atom; the energy of that drop equals the energy of a 589 nm photon, found with E = hc/lambda. If a sample contains two elements, you simply see both sets of lines superimposed, undiminished, since each element's electrons transition independently of the other's.

Flip the setup around and shine white light through a cool gas instead of heating it, and you get an absorption spectrum: a continuous rainbow with dark lines missing at those same characteristic wavelengths, because ground-state electrons absorb exactly those photons to jump up. Emission and absorption are two views of the identical transitions, one showing photons leaving, one showing photons being removed from a beam. This same logic of matching line patterns to known references threads through other labs in this unit, like identifying unknowns from their properties or verifying a solution's identity.