Light and Matter
Scientists and philosophers have argued for centuries about the nature of light and matter. As far back as 400
B.C. philosophers have debated about the structure of matter. The Greek philosopher Democritus (c. 460 – c. 370
BC) believed that matter “consists of an infinite of minute particles”. Later, the famous Greek scientist and philosopher Aristotle (384 – 322 BC) believed matter was continuous, or that it could be broken down infinitely. It was Aristotle’s influential logic on this topic (and many others) that was to be felt for the next 2000 years.
Centuries later Englishman Isaac Newton (1642-1727) stated that light was a stream of particles. A contemporary
Dutch physicist Christian Huygens (1629-1695) believed light was a wave. This issue was resolved by James Clerk
Maxwell (1831-1879) of Scotland who showed the properties of light could be explained by an electromagnetic wave. Maxwell showed the electric and magnetic field of light moved perpendicular to each other and perpendicular to the direction of the wave. All waves could be described in terms of velocity, v; frequency, f and wavelength,  . The wave equation is v = f  . The speed of all electromagnetic waves (light, UV, radio) is 3 x 10 8 m/s.
Maxwell’s wave theory had to be changed about a hundred years ago when discoveries were made when light interacted with matter and had an effect that puzzled scientists. The photoelectric effect was discovered by
German Heinrich Hertz in 1888. When light, particularly light of high energy shines on the surface of a metal, electrons are emitted from the surface. In 1900, Max Planck (1858-1947) discovered different metals had different thresholds when they released electrons. This threshold was based on wavelength/ frequency of the light not the intensity. Planck suggested that matter, at the atomic level, can absorb or emit only discrete quantities of energy. Each of these quantities is called a quantum of energy. In essence, Planck is saying that energy is not continuous but comes in packages. Analogy: in the past scientists thought of energy in terms of a ramp and now think if energy like a ladder.
In 1905, an unknown Swiss clerk named Albert Einstein published three papers. The first paper gave his explanation of the photoelectric effect. He proposed light was made up of small bundles of light energy, called quanta or photons and that the E = hf (h is known as Planck’s constant). The second paper explained Brownian motion (gas behaviour). The third paper was about the now famous link between matter and energy, E = mc 2 .
In 1913, Danish physicist Niels Bohr gave his theory on quantized energy levels that he believed existed in all atoms. When heat or electrical energy excited the electrons they rose to higher energy levels. Upon falling back down to the ground state, they emit light energy of varying wavelengths. Through his work with New Zealand’s
Ernest Rutherford, Bohr was able to ascertain energy levels, n; and the maximum number of electrons at each energy level 2n 2 .
In 1924, French physicist Louis de Broglie suggested that since light has wave/particle properties so might elementary particles (p + , n o , e ). The idea that matter may be represented by a wave was a new idea and gave rise to the equation:  = h /mf
In 1926, after Austrian Erwin Schrodinger heard of de Broglie’s “electron wave” it occurred to him that this idea could be used to solve the electron location problem in an atom. Schrodinger proposed using wave (quantum) mechanics – a new branch of mathematics - and applying it to the behaviour of the electron. The result was our present orbital theory of the electron. Orbitals, electron clouds, take on certain shapes or sublevels (s-spherical, p-hourglass, d-cloverleaf). The energy level, n, must equal the number of sublevels. Within a sublevel the number of orbitals will change. For the s – sublevel there is 1 orientation, for the p-sublevel, 3; for the d-sublevel, 5.
At the same time, Swiss physicist Wolfgang Pauli proposed why electrons do not crowd around the lowest (1s) level.
The electron is distributed around the nucleus as a cloud of probability and they were known to spin and have an electric field. The Pauli Exclusion Principle states: No orbital can accommodate more than two electrons and in order for electrons to exist together in the same orbital, they must be opposite in spin.
In 1927, German Werner Heisenberg concluded there are limitations on which an electron can be located. The
Heisenberg Uncertainty Principle states: There is an uncertainty in determining the position and momentum of an electron. If one is known more accurately the other is more uncertain. In 1932, James Chadwick working with
Rutherford, discovers the neutron thus explaining the existence of isotopes.