Views of fundamental composition of matter have changed over many centuries.
Beginning with Democritus in ancient Greece, it was thought that the world
was made of small "atoms" (Greek for "indestructible") that were indivisible
any further. Then in the late 1800's J. J. Thompson, upon the discovery
of the electron, conjectured that the "atom" was made of negatively-charged
particles (electrons) dispersed among positively-charged particles (protons)
as a "plum pudding".
Rutherford's model
Thompson's model
Bohr's model
Thompson's model was modified by Ernest Rutherford to explain gold-foil
experiments. Rutherford's model consisted of a central collection of protons
with electrons in "planetary orbits" about the nucleus. The problems with this
model were:
1. Atomic spectra was still unexplained. It was widely known that elements,
upon heating, would give off radiation, found to be at discrete frequencies
after
passing through a spectroscope. The spectral lines were unique
to each element.
2. Rutherford's solar-system model could not maintain stability. Since the
electrons were constantly changing direction, they were accelerating, which
would continually radiate energy. The electrons would eventually spiral into
the nucleus and the atom would collapse.
Then Niels Bohr in 1913 introduced his Bohr Atom. Based on the quantum theory, it proved to be more successful than
Rutherford's
picture at explaining atomic properties. Bohr proposed that electrons orbited
the nucleus, but the electrons contained enough energy to match the electric
pull
of the protons. This way, the atomic stability would be preserved. He also
said that electrons could occupy only certain orbits, or energy-levels.
As Planck suggested, the energies of the electrons had to be integral multiples
of a particular quantum of energy. When an atom absorbed a photon (a packet of
electromagnetic radiation), an electron would jump up orbitals corresponding to
the amount of energy held by the photon. When an electron returned to lower
energy
levels, it emits a photon with energy equivalent to the difference between the
two levels.
When a photon has a certain energy, it has a certain frequency. By Planck's
equation:
E = h * f
the higher the frequency, the higher the energy. Differing properties of energy
levels of different elements would cause the frequencies
of atomic spectra to be unique to each element. Thus Bohr's atom was better at
explaining electron orbits and spectra.
However well the Bohr's atom fit the experimental data, scientists did not know
exactly
why only orbits at certain energies were allowed. In trying to explain the
problem,
Louis De Broglie suggested that electrons were waves.
When the electron-wave did not orbit back on to itself, the wave would cancel
itself out,
and the electron would phase out of existence. Hence, the electron could exist
only
when the wave was at harmonic distances from the nucleus.
Though De Broglie's idea was not entirely physically accurate, the introduction
of
waves paved the way for quantum mechanics to further develop.
