Bohr

In 1900 Max Planck was attempting to develop an equation to match the energy emitted from an ideal emitter called a black body.  (A kiln emits energy through a small hole as a black body.)  Encountering difficulties, Planck tried a calculus trick of considering the energy emitted in small discrete amounts.  Anticipated the size of the small energy lump would be irrelevant, Planck realized the radical results when he found the size of the lump was not irrelevant.  The lump size is proportional the frequency of the radiation times a tiny number now called Planck's constant = 6.6 x 10^-34 Joule-seconds.  Planck thought the lumpy (quantum) emission of light energy might be due to the lumpy nature of atomic matter.

In 1905 Albert Einstein recognized that Planck's constant in another equation describing the photoelectric effect, E = hν - W, implies that the term  hν  requires light by its nature to be quantized.

Rutherford created a dilemma when he proposed that each atom has a nucleus containing nearly all the atom's mass and a positive charge equal the atomic number.  If an atom's negative electrons were stationed in the remaining atomic space, the electrical attraction would pull the electrons toward the nucleus, collapsing the atom.  Alternatively, if the electrons were placed in motion orbiting the nucleus, the electromagnetic laws governing accelerated electric charges would require that each electron broadcast away its energy resulting in a spiral crash into the nucleus.  Neither stationary nor moving electrons could form stable atoms.  In 1913, Niels Bohr (1885-1962 at left), after working with Rutherford, proposed that electrons did orbit the nucleus but that the electromagnetic laws somehow acted differently inside atoms.  Bohr proposed that since Planck's constant had the units of angular momentum, perhaps the angular momentum of each electron was specified by an integer (quanta) multiplied of Planck's constant.  Since electrons would not be permitted to have fractional quanta, an electron could not gradually broadcast.

But Bohr proposed that an electron could absorb the correctly sized quanta of energy to be excited to any higher orbit allowed by a larger integer.  Then the excited electron might spontaneously returning to a lower orbit, emitting the discrete color of light specified by the difference in energy of the two electron orbits.  Bohr calculated the orbits possible, the energies a electron would have in each orbit, and the spectra colors for each orbit transition in the simplest atom, hydrogen.  He found a perfect match with the spectra of hydrogen known at the Lyman, Balmer, Paschen, Brackett and Pfund series.

Noting that the periodic table has inert elements periodically, with elements that easily lose electrons on one side, and atoms that take electrons on the other side, Bohr proposed that filling or emptying an orbit that holds a specified maximum number of electrons could explain the table.  This proposal made it possible to predict where other missing elements might be discovered.

G.N. Lewis, Walther Kossel, and Irwing Langmuir in 1916 developed a new theory of chemical bonding related to full electron shells.  Ions in polar materials such as salts gain or lose electrons so that the outer electron shell is full.  Ions then bond by electrical attraction.  Non-polar molecules such as in organic molecules share electrons in a covalent bond that fills the outer electron shells.  Orbits may contain a limited number of electrons.  Inert gases are atoms with full electron orbits.

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created 23 March 2002 latest revision 4 November 2006

by D Trapp