Quantum mechanics attempts to
describe and account for the properties of molecules and atoms and their
constituents—electrons, protons, neutrons, and other more esoteric particles
such as quarks and gluons. These properties include the interactions of the
particles with one another and with electromagnetic radiation (i.e., light, X
rays, and gamma rays).One of the most remarkable scientific advances of the 20th century is the development of quantum mechanics - the description of the behavior of matter on the atomic and sub-atomic scale. It is now a powerful tool for understanding the behavior of atoms and molecules, and is vital to physicists, chemists and biochemists alike.
Its roots lie in Max Planck's discovery at the turn of the century that the radiation from a hot object can be successfully described only if it occurs with specific amounts of energy - "quanta" - rather than with a continuous range of energies. This discovery led ultimately to the description of light in terms of "particles", known as photons, the name coined in 1926 by the American Gilbert Lewis.
In 1913 the Dane Niels Bohr built on these ideas to postulate that the energy of the atomic electrons must also be "quantized". The model explained the origin of the spectra of light emitted by atoms such as hydrogen, which had long been recognized to have characteristic and separated lines of color. But the explanation of why the energy of the electrons should be quantized had to wait until the mid-1920s with the full development of the mathematical formulation known as quantum mechanics by the Austrian Erwin Schrodinger, the German Werner Heisenberg and the British physicist Paul Dirac.
Schrodinger's theory of quantum wave mechanics treated the electron with a wavelike description, the amplitude of the wave giving the probability of finding the electron at a given point in space and time. This wave, like the electromagnetic waves of radiation, was subject to quantization, and the energy levels (shells) in Bohr's model could be explained in terms of the allowed energies of an electron-wave, effectively caught by the electric attraction of the nucleus.
One of the most fascinating discoveries to emerge
from quantum theory is the "uncertainty principle" found by Werner
Heisenberg. This tells us that it is impossible to measure both members of certain
pairs of properties to arbitrarily great accuracy. The better one is known, the
worse becomes our knowledge of the other, rather as the illustration background
in a photograph is blurred if we pan the camera to catch a sharp picture of a moving
object. Thus if we try to pin down the position of an electron, we lose our knowledge
of its momentum. This is because, at the subatomic level, even "looking"
at an electron requires a photon of light, and this alters the electron's energy.
Indeed, the closer we try to look, with light at shorter wavelengths, the more
energy we impart to the electron. (Adapted from ‘Science A History of Discovery
in the Twentieth Century’, by Trevor Williams, and ‘Encyclopedia Britannica’)
