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’)