Photons, Electrons, and Atoms
Visible and non-visable light Frequencies around Hz Much higher than electric circuits Theory was about vibrating charge in the atom This lead to the “quantum” theory of light It states that some of the properties of EM radiation resemble those of particles Where KE is proportional to f
The Photoelectric Effect It is caused by an element absorbing certain wavelength of light which causes electrons to be ejected from the surface This “potential energy barrier” can be found with the work function Classic mechanics predicted that the amplitude of the wave, not its frequency, would cause more electrons to be emitted
Photoelectric cont. It was found that the amplitude, no matter how bright, would not cause the electrons to leave. Only when critical or threshold frequency was reached, would the electrons move For most metals it’s in the ultraviolet range At frequencies higher than ft, the stopping potential would completely stop the flow of electrons from the surface
Equations Stopping potential Energy of a photon Planck’s constant h = x j.s
The eV The electron volts relates electrons and energy 1 eV = x J Ex: If V o = 5V then e V = 5eV Example 41-1 Photon energy p = momentum
EM Spectrum
Line spectra Line spectra phenomenon vs spectrum Each line is an image of the spectrograph slit, deviated at an angle that depends on the frequency of the light Each element emits a “signature” Certain wavelengths of light unique to its atomic structure Classical mechanics could not explain this
Energy Bohr proposed the idea of energy levels All atoms of an element have the same energy levels, with no electrons existing in “intermediate” levels The electrons can “jump” from a higher to a lower level by emitting a photon hf = E i - E f
Balmer Series Trial and error resulted in the formula R = x 10 7 m -1 Substituting in for we get En ( energy levels
Emission spectra
Sources s/hillchem3/medialib/media_portfolio/07.ht mlhttp://cwx.prenhall.com/bookbind/pubbook s/hillchem3/medialib/media_portfolio/07.ht ml