In the late 18th century, many physicists believed that they had made great progress in physics, and there wasn't much more that needed to be discovered. The classical physics at the time was widely accepted in the scientific community. However, by the early 20th century, physicists discovered that the laws of classical mechanics break down in the atomic world, and experiments such as the photoelectric effect completely contradict the laws of classical physics. As a result of these crises, physicists began to construct new laws of physics which would apply to the atomic world; these theories would be collectively known as quantum mechanics. Quantum mechanics, in some ways, completely changed the way physicists viewed the universe, and it also marked the end of the idea of a clockwork universe (the idea that universe was predictable).
Electromagnetic radiation (ER) is a form of energy that sometimes acts like a wave, and other times acts like a particle.
Visible light is a well-known example.
All forms of ER have two inversely proportional properties: wavelength and frequency.
Wavelength is the distance from one wave peak to the next, which can be measured in meters.
Frequency is the number of waves that pass by a given point each second.
Since wavelength and frequency are inversely related, their product (multiplication) always equals a constant—specifically, 3.0 x 108 m/sec, which is better known as the speed of light.
The wavelength and frequency of any specific occurrence of ER determine its position on the electromagnetic spectrum.
Again, the conversion of the wavelength of a light is determined by using,
where c is the constant 3.0 x 108 m/sec, the speed of light in a vacuum, λ = wavelength in meters and v=frequency in hertz, 1/s. It is important to note that using this equation, one can determine the wavelength of light in frequency.
The Discovery of the Quantum
So far we have only discussed the wave characteristics of energy. However, the wave model cannot account for something known as the photoelectric effect. This effect is observed when light focused on certain metals apparently causes electrons to be emitted. (For a more comprehensive discussion of the photoelectric effect, see the associated atom in this module. ) For each metal, there is a minimum threshold frequency of electromagnetic radiation that is needed to be shone on it in order for it to emit electrons. One could not replace a certain amount of light at one frequency with twice as much light of half the frequency. If light only acts as a wave, the effect of light should be cumulative—the light should add up, little by little, until it causes electrons to be emitted. Instead, there is a clear-cut minimum of the frequency of light that triggers the electron emissions. The implication of this is that frequency is directly linked to energy, with the higher light frequencies having more energy. This observation led to the discovery of the minimum amount of energy that could be gained or lost by an atom. Max Planck named this minimum amount the "quantum," plural "quanta," meaning "how much. " One photon of light carries exactly one quantum of energy.
Planck is considered the father of the Quantum Theory. According to Planck, each energy element E is proportional to its frequency ν: E=hv, where h is Planck's constant. Planck (cautiously) insisted that this was simply an aspect of the processes of absorption and emission of radiation, and had nothing to do with the physical reality of the radiation itself. However, in 1905 Albert Einstein interpreted Planck's quantum hypothesis realistically and used it to explain the photoelectric effect, in which shining light on certain materials can eject electrons from the material.
More Evidence for a Particle Theory of Energy
When an electric current is passed through a gas, some of the gas molecules' electrons move from their ground state to an excited state that is further away from their nuclei. When the electrons return to the ground state, they emit energy of various wavelengths. A prism can be used to separate the wavelengths, making them easy to identify. If light acted only as a wave, then there should have been a continuous rainbow created by the prism. Instead, there were discrete lines created by different wavelengths. This is because electrons release specific wavelengths of light when moving from an excited state to a ground state.