http://www.ck12.org Chapter 23. The Atom
Problems with the Planetary Model
The planetary model had two major problems immediately. According to classical mechanics, a charged particle
being accelerated (speeding up, slowing down, or turning a corner) would always emit electromagnetic radiation.
That means that an orbiting electron would constantly emit energy and thus move closer to the nucleus until
eventually it would collapse into the nucleus. Clearly this doesn’t happen and thus it opposes the idea that electrons
move like planets. The second problem was that atoms gain energy (in the form of heat or light) and re-emit the
energy in an exact set of light frequencies. The magnitude of the energy involved in this light emission is too small to
be involved with nuclear changes and therefore, the electron configuration must be responsible for the light emission
by atoms. The planetary model offered no explanation for the light spectrum of atoms.
A major clue to the electron arrangement in an atom came from studying the light emitted by atoms. When electricity
is passed through gaseous atoms, the atoms emit a spectrum of light that is specific for that element. The hydrogen
atom, for example, emits a pinkish light but when that light is passed through a prism, we see a very few frequencies
of light that are quite specific for hydrogen. The energies of these light frequencies is much too small to be involved in
the nucleus of atoms, therefore, any explanation of these wavelengths would have to involve the electron arrangement
in atoms.
The Bohr Model of the Atom
Niels Bohr attempted to join the nuclear model of the atom with Einstein’s quantum theory of light and with his own
idea of electron energy levels to explain the electron arrangement within the atom. Bohr started with the planetary
model of electron arrangement but postulated that electrons in stable orbits would not radiate energy even though the
electrons were being accelerated by traveling in circular paths. Bohr hypothesized that the electrons were organized
into stepwise energy levels within the electron cloud and only radiated energy when the electrons moved from one
energy level to another. Bohr’s hypothesis suggested that the energy of atomic electrons came in packages and only
whole packages could be absorbed or emitted. This quantization of energy allowed electrons to only absorb or emit
exact amounts of energy to move from one energy level to another.
The quantization of energy is not apparent in everyday experience. If we could observe molecular sized automobiles
traveling down miniature roads, we would see cars traveling at 7 miles per hour, or 14 miles per hour, or 21 miles
per hour, but never at 9 or 17 miles per hour. The quantization of energy means that energy comes in packages and
when energy is added to an object, whole packets of energy must be added. This is the explanation for why atomic
electrons are only allowed to have certain amounts of energy and therefore, occupy certain energy levels. The lowest
energy level for an electron is near the nucleus and each quanta of energy added moves the electron to the next
distant energy level. Einstein’s theory says that each light photon has an energy ofh f, wherehis Planck’s constant,