General Chemistry/Introduction to Quantum Theory
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[edit] Introduction to Quantum Mechanics
In the late 18th century, many physicists believed that they have made great progress in physics, and there wasn't much more that needed to be discovered. The classical physics at the time was very well widely accepted in the scientific community. However, by the time of 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 with 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 clockwise universe (the idea that universe was predictable).
[edit] Electromagnetic Radiation
Electromagnetic radiation, such as light, is a form of energy that sometimes acts like a wave, and other times acts like a particle. Visible light is the most well known example. Different forms of ER have two inversely proportional properties: wavelength and frequency. Wavelength is the distance from one peak to the next, which can be measured in meters. Frequency is the number of waves that pass through a certain point per 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 ER determines its position on the electromagnetic spectrum.
As you can see, visible light is only a tiny fraction of the spectrum.
The energy of an electromagnetic wave is given by E = hf, where h is a constant and f is the frequency. Energy is directly proportionate to frequency—doubling the frequency will double the energy.
[edit] The Discovery of the Quantum
So far we have only discussed the wave characteristics of energy. However, waves cannot describe something known as the photoelectric effect. This is when, if you shine light on some metals, they emit electrons. (This is how the "photoelectric" solar panels work.)
For each metal it was found that there was a minimum frequency of electromagnetic radiation that needed to be shone on it in order for it to emit electrons. You could not replace a certain amount of light at one frequency with twice as much of other light that was half the frequency. If light only acted as a wave, the effect of light should have been cumulative - the light would add up, little by little. Instead, there was a clear-cut minimum of the frequency of light that could be used.
The implication of this was that frequency is directly linked to energy, 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.
[edit] More Evidence for a Particle Theory of Energy
When an electric current is passed through a gas, some of its electrons move from their ground state to an excited state that is further away from the nucleus. When the electrons return to the ground state, they emit energy of different 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 definite lines created by different wavelengths. This is because electrons release specific wavelengths of light when moving from an excited state to a ground state.
