Title ~ Physical Properties of Gases, Liquids and Solids
Author ~ R. P. W. Scott
Section ~ The Bohr Atom.
The Bohr Atom
The
Bohr atom is depicted in figure 1 as a hydrogen atom. Although now considered
obsolete, the Bohr atom serves as a useful introduction to the structure of the
atom and, thus, will briefly be discussed. Bohr depicted the hydrogen atom as a
central positively charged nucleus (in the case of hydrogen, a single proton)
orbited by a negatively charged electron held in equilibrium by the balanced
outward centrifugal force and the inward electrical attractive force. This
likens the electron and atom to a planet revolving round a star.
Figure 1. The Bohr Atom
The
electron could spin round a number of different orbits having different
energies in each, the actual energy increasing as the orbit becomes larger.
Three such orbits are depicted in figure 6, labelled (n=1), (n=2), and (n=3).
The innermost orbit has the lowest energy (n=1).
If
light of 656 nm wavelength strikes a hydrogen atom and is absorbed (
) and this can result in an
electron situated in the orbit where n=2 being transferred to the orbit where
n=3. In a similar manner if an electron in orbit n=3 falls back to the obit
where n=2 then light of 656 nm will be emitted. Thus, the basic mechanics of
UV/visible light absorption and fluorescent emission can be accounted for.
However, although the energy difference of the electrons between orbits could
be understood, an explanation of the factors that control the value of (n)
requires a more sophisticated model to be considered. The star-planet model
needed to be significantly modified.
In
1924 the concept was introduced that all electromagnetic radiation could be
considered as either waves or particles and furthermore very small particles (i.e. electrons) travelling at high speed
could also exhibit wave properties. This was confirmed by experiments that
demonstrated the diffraction of electrons (in the manner of light) and
ultimately the development of the electron microscope.
In
1924, Broglie produced equations that reconcile, in a
relatively simply way, this particle-wave nature as follows.
If
(
) is the wavelength of the wave and
(p) the momentum of the particle,
Then
Where
h is Plank’s constant.
And
for an electron, mass (me)
travelling in a straight line at a velocity (v)
Then p = mev
The wave nature of the electron easily explains the restricted
nature of the different orbits. The
electron can only exist in orbits in which the wave is a standing wave.
Such
a condition is depicted in figure 2. It follows that if the orbit radius is (r), for a standing wave the circumference of the
orbit (
) must be an integral number of wavelengths, i.e., (
).
The integer (n) is known as the principal quantum number introduced by Bohr and can
take values of 1.2.3………¥
Figure 2. An Electron in Orbit Presented as a
Standing Wave
The
orbital depicted in figure 2 is for n =12.
The
behaviour of an electron as a wave rather than as a particle provoked a quite
different approach to the theory called Quantum Mechanics or Wave mechanics.
The subject of wave mechanics is outside the scope of this book and for those
readers who wish to study the subject further they are recommended to read
Basic Atomic and Molecular Spectroscopy by J. Michael Hollas
published by the Royal Society of Chemistry.
The physicist, however, may split the atom into electrons and nuclei, then the nuclei into protons and neutrons and with use of very high electrical or magnetic energy the protons and neutrons into quarks, gluons etc. With the start of the twentieth century the earth was no longer thought to consist of the four elements, earth, air, fire and water’ but of ‘ether, matter, and positive and negative electricity’. This was coupled with the hopeful theory that matter might eventually be shown to be comprised of positive and negative electricity. The introduction of the concept of ‘ether’ needs some discussion.