MPS - Electron Energy Levels
Electron Energy Levels
Neils Bohr envisioned a model of the atom where electrons orbit the nucleus of the atom at discrete energy levels. He referred to these orbits as stationary states. What this means is that an electron must gain or lose a quantized amount of energy to move between states. To move up an energy level, an electron must gain the difference in energy between its current state and the higher, excited state. This is done by absorbing a photon with that energy. Each level is represented by a quantum number. The lowest energy level has a quantum number of 1 and is known as the ground state (E1). The higher states are labeled E2, E3, etc. and are known as excited states. When an electron falls from an excited state, it emits a photon with the exact energy as the difference between the states. Remember that a photons energy is directly related to a frequency by the equation E = hf. Different elements have different values for the energies of the various electron states. This means that, when falling from an excited state, different elements will produce different frequencies of light. This is the physical basis for the atomic spectra you studied in chemistry.
A hydrogen atom produces one of the simplest emission spectra. These visible lines represent specific frequencies of light as seen through a diffraction grating (used to separate the different frequencies). Each frequency corresponds to the difference in specific energy levels. The visible spectral lines, called the Balmer series after J.J. Balmer, correspond to the differences in energies between the quantum states n = 2 & n = 3, n=2 & n=4, and n=2 & n=5. To calculate the frequencies of photon necessary to excite an electron to any higher state we simply use the equation E=hf, but substitute the difference in state energies, 
for the photons energy.
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