The most important properties of atomic and molecular structure may be
exemplified using a simplified picture of an atom that is called the
Bohr
Model. This model was proposed by Niels Bohr in 1915; it
is not completely correct, but
it has many features that are approximately correct
and it is sufficient for much of our discussion. The correct theory of the
atom is called
quantum mechanics; the Bohr Model is an approximation
to quantum mechanics that has the virtue of being much simpler.
(Here is a
more realistic discussion of what atomic orbitals look like in quantum
mechanics.)
A Planetary Model of the Atom
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The Bohr atom
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The Bohr Model is probably familar as the "planetary model" of the atom
illustrated in the adjacent figure that,
for example, is used as a symbol for atomic energy (a bit of a misnomer, since
the energy in "atomic energy" is actually the energy of the nucleus, rather
than the entire atom). In the Bohr Model the neutrons and protons
(symbolized by red and blue balls in the adjacent image) occupy a
dense central region called the nucleus, and the electrons orbit the nucleus
much like planets orbiting the Sun (but the orbits are not confined to a plane
as is approximately true in the Solar System). The adjacent image is not to
scale since in the realistic case
the radius of the nucleus is about 100,000 times smaller than the
radius of the entire atom, and as far as we can tell electrons are point
particles without a physical extent.
This similarity between a planetary model and the Bohr Model of the atom
ultimately arises because the attractive
gravitational force in a solar system
and the attractive
Coulomb (electrical) force between the positively charged nucleus and
the negatively charged electrons in an atom are mathematically of the same form.
(The
form is the same, but the intrinsic
strength of the Coulomb
interaction is much larger than that of the gravitational interaction; in
addition, there are positive and negative electrical charges so the Coulomb
interaction can be either attractive or repulsive, but gravitation is always
attractive in our present Universe.)
But the Orbits Are Quantized
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Quantized energy levels in hydrogen
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The basic feature of quantum mechanics that is incorporated in the Bohr Model
and that is completely different from the analogous planetary model is that the
energy of the particles in the Bohr atom is restricted to certain discrete
values. One says that the energy is
quantized. This means that only
certain orbits with certain radii are allowed; orbits in between simply don't
exist.
The adjacent figure shows such quantized energy levels for the hydrogen atom.
These levels are labeled by an integer
n
that is called a
quantum number.
The lowest energy state is generally termed the
ground state. The states
with successively more energy than the ground state are called the
first excited
state, the
second excited state, and so on. Beyond an energy
called the
ionization potential
the single electron of the hydrogen atom is no
longer bound to the atom. Then the energy levels form a continuum. In the case of
hydrogen, this continuum starts at 13.6 eV above the ground state ("eV" stands
for "electron-Volt", a common unit of energy in atomic physics).
Although this behavior may seem strange to our minds that are trained
from birth by
watching
phenomena in the macroscopic world, this is the way things behave in the
strange world of the quantum that holds sway at the atomic level.
Atomic Excitation and De-excitation
Atoms can make transitions between the orbits allowed by quantum mechanics by
absorbing or emitting exactly the energy difference between the orbits. The
following figure shows an atomic excitation cause by absorption of a photon and
an atomic de-excitation caused by emission of a photon.
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Excitation by absorption of light and de-excitation by emission of light
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In each case the
wavelength of the emitted or absorbed light is exactly such that the photon
carries the energy difference between the two orbits. This
energy may be calculated by
dividing the product of the Planck constant and the speed of light
hc by the wavelength of the light). Thus, an atom can absorb
or emit only certain discrete wavelengths (or equivalently, frequencies or
energies).