Geoscience Reference
In-Depth Information
In 1903, just before Einstein's theory of unity, Thomson postulated that individual atoms
were spheres of “uniform positive electrification,” scattered with electrons rather like “cur-
rants” in a bun. Thomson also realized that because many atoms appeared to be electrically
neutral, other “positively charged” subatomic particles must also exist within the atom.
Shortly after this proposed model of Thomson, in 1910 Lord Ernest Rutherford and his
researchers led to the proposition that an atoms mass must be concentrated at its center,
that is,. the nucleus (Rutherford,
1911
). Much of Rutherford's work was complemented
by the Danish physicist Neils Bohr, who, in
1913
, proposed that electrons exist in
quan-
tized
states. Bohr's physical model postulated that the energy of these quantized states was
determined by the angular momentum (motion through space) of the electron's orbit about
the nucleus. Quantized states do not vary “continuously” but rather in permitted quantum
leaps, that is, between precise values. Furthermore, electrons were free to leap between
these states, or orbits, by the emission or absorption of photons at discrete frequencies.
Bohr used the notion of quantized orbits to account for the emitted spectral lines of hydro-
gen atoms. Although momentous in our understanding of physics, Bohr's model failed to
predict the observed relative intensities of spectral lines, and more importantly the spectra
of more complex atoms with fine and hyperfine structure. Despite the shortfalls of Bohr's
theory, which was constrained to the simplest known atom, hydrogen, the notion that an
atom is a dense nucleus of positive charge surrounded by lower-mass orbiting electrons
was an established idea by 1914.
Bohr's initial model (Bohr,
1922
) helped scientists advance our understanding of chem-
ical bonding between atoms and better understand the quantum state. In 1916, American
scientist Gilbert Newton Lewis proposed the idea of the covalent chemical bond, in which
the bond between two atoms is maintained by a pair of “shared” electrons. The work of
Lewis was elaborated further in 1919 by the American chemist Irving Langmuir. Langmuir
suggested that all electrons were distributed in consecutive spherical “shells” of equal
thickness. Langmuir further divided these shells into a number of cells each containing one
pair of electrons. Using this model Langmuir was able to explain the chemical properties of
all elements in the periodic table according to the periodic law, which states that the chem-
ical properties of the elements are periodic functions of their atomic numbers.
In 1923, Walter Heitler and Fitz London fully explained electron-pair formation and
chemical bonding in terms of quantum mechanics (Heitler and London,
1927
). In the same
year the French physicist Louise de Broglie proposed that wave-particle duality applied
not only to photons, but also to electrons and every other subatomic physical system; this
work was published in his PhD thesis in 1924. Austrian physicist Wolfgang Pauli (
1925
)
observed that the shell-like structure of the atom could be explained by a set of four param-
eters that define every quantum energy state, as long as each state was inhabited by no more
than a single electron.
These parameters are:
•
Principle quantum number,
n
. In Bohr's model this number largely determines the energy
level and the average distance of an electron from the nucleus.
