Leading up to the Quantum Theory

 exhibits wavelike behavior  moves at a speed 3.8 × 10 8 m/s in a vacuum  there are measureable properties of light (wavelength and frequency) and these properties are related to each other mathematically c = λν

 No electrons were emitted if the light’s frequency was below a certain minimum, regardless of the intensity of the light. If light is thought of only in a wave context, then light of any frequency should supply enough energy to eject an electron. This was the conundrum scientists were trying to explain.

 Max Planck – suggested that energy comes in small packets called quanta. He proposed the relationship between a quantum of energy and radiation to be E = hν  E=the energy of a quantum of radiation; ν=frequency of the radiation emitted and; h=constant (now a fundamental in physics called Planck’s constant = 6.626× Js)

 Albert Einstein expanded on Planck’s Theory and introduced an idea that all electromagnetic radiation has a wave and particle nature to it. Each particle of light carries a quantum of energy. These particles are called photons. E photon = hν

 Thus Einstein explained the photoelectric effect by proposing that radiation is absorbed by matter in whole numbers of photons. In order for an electron to be ejected, the electron must be struck by a single photon possessing a minimum amount of energy (which corresponds to the frequency).

Classical theory of energy predicted that hydrogen atoms would be excited by whatever amount of energy added to them – then you would expect to see a continuous spectra of frequencies.

What we really saw was four very specific lines at 656 nm (red), 486 nm (green), 434 nm (blue) and 410 nm (purple). Using the photon idea…These specific frequencies indicated that the energy emitted when an electron falls from an excited state to its ground state or a lower energy excited state. This also indicated that an atom’s energy states were fixed.

 Proposed a model that linked the atom’s electron emission to photon emission. Electrons can only circle the nucleus only in allowed paths or orbits; the further the obit, the higher the energy.  His model could explain emission spectrum observed for the hydrogen atom. He calculated the allowed energy levels for the hydrogen atom and related the possible energy level changes to the lines in the emission line spectrum. His calculated values agreed with experimentally observed values.  What his model didn’t do: (1) Did not explain the spectra of atoms with more than one electron. (2) Did not explain the chemical behavior of atoms.

 de Broglie wondered that if light has a wave- particle duality, then maybe electrons may have the same nature.  Scientists knew any wave confined to a space can have only certain frequencies. De Broglie suggested that electrons be considered waves confined to the space around an atomic nucleus.  Experiments demonstrated that electrons, like light waves, can be bent or diffracted (bending as it passes through a small opening). Also, some experiments showed that electron beams interfere with each other.

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 Heisenburg worked on answering the question of “where” electrons are in the atom, if they (the electrons) are both wave and particle.  Electrons are detected by their interactions with photons. Because photons have the same energy as electrons, any attempt to locate them with a photon knocks the electron off its course.  The Heisenburg Uncertainty Principle states that it is impossible to determining both the position and velocity of an electron at the same time.

 Schrödinger used the dual wave-particle nature to create an equation that treated electrons as waves. With this equation and Heisenburg’s Uncertainty Principle laid the foundation for our Modern Quantum Theory of the atom.  Rather than electrons travelling in distinct orbits around the nucleus, electrons occupy three- dimensional regions of high probability around the nucleus.

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