|CH301H - Principles
of Chemistry I: Honors
Fall 2011, Unique 51040
Lecture Summary, 22 September 2011
The electromagnetic spectrum can be resolved into components of
vastly different frequency and wavelength. Although the visible
spectrum covers only a tiny fraction of the electromagnetic spectrum
(about 2.5 eV), because our eyes are sensitive to only this spectrum,
we spend a lot of time thinking about how matter interacts with light
at these wavelengths. By discharing an electrode into a pure gas,
collisions between the free electrons and the molecules of gas will
provide energy to move electrons in the gas to higher energy levels.
When that electron falls back (or "relaxes") to its lowest energy
(or "ground") state, it will emit energy in the form of light. We
looked at a few examples of noble metel gasses that emit light in the
visible region of the spectrum, and saw that sometimes the gas will
emit several different wavelengths of light. You can perform this
experiment on your own using a PhET Java applet:
By the end of the 19th century, these spectra lines had been measured for many atoms, and it was known empirically that the energy (i.e. frequency) of light emitted in these experiments was quantized.
Bohr Model of the Atom: In 1913, Bohr attempted to define a model to explain the stable atom. He started with two assumptions: 1) an electron orbits around a positively charged nucleus according to Rutherford's planetary model, and 2) the angular momentum of the orbiting electron is quantized. From these two postulates, he derived an expression for the energy of a one-electron atom based only on constants and the "quantum number" n. The derivation of this expression is given here:
The Bohr model was able to explain the ionization energies and observed emission spectra of one-electron atoms.