The Origin of Light
|The example on this page looks at a Hydrogen atom that begins in the energy state of n=3 (Don't worry, it is just a system used to count the energy levels.) Notice that at n=3 there are 3 complete waves in that loop that is the electron.|
|Creating more waves in the same length of a spring requires more energy, (you might remember this from the spring lab). If more energy was added, it may or may not move the electron to the higher energy level. If it is only enough energy to take it to n=3.25, the wave will stay at the n=3 level. This is because the end of the wave and the beginning of the wave must be able to connect in phase. At n=3.25, one end of the wave is on its way up, the other on its way down. Physically, they can't connect. For this reason we would say that n=3.5 or 3.8 or 2.9 or 11.6 is not "allowed". In fact, only whole numbers of waves are allowed. We say that the energy levels are "quantized" to certain amounts.||
|If just the right amount of energy has been put in, then the electron will jump to being a wave with 4 complete waves.|
Then a photon (wave packet) of light comes in from the left. The energy of the photon just happens to be exactly the amount needed to raise the electron up to the next energy level of n=4. Notice at n=4 there are 4 complete waves. We would say that an electron at n=4 is more excited or energetic than an electron at n=3. Notice that the frequency of the waves are higher at n=4.
Due to the desire to keep lower energy levels full, the electron eventually drops back down to n=3 and ejects a photon of light. This photon is exactly the same energy (same color) as the photon it absorbed. This leads to the idea that an atom can only produce the colors of light (emission spectrum) that match the colors they can absorb (absorption spectrum).
There are three basic ways to excite an electron:
- Collide it with another atom or electron.
- Heat it up. (thermal excitation)
- Bombard it with a photon of just the right energy (color).