Download Spectra and atomic structure

SPECTRA OF THE ELEMENTS
So far we have thought about the radiation emitted by hot solids but we are now going to broaden it
by considering gases. Looking at light emitted by a source through a prism or a diffraction grating
can tell us a great deal, not only what the substance is but also much about its atomic structure!
Emission spectra
You should have a look at various spectra (some of which are shown in the photographs). You will
see that they fall into roughly three groups.
(i) Line spectra - these are emitted from hot monatomic gases - the atoms are not linked to each
other in any way.
(ii) Band spectra - these are emitted by gases as well but from ones with more than one atom per
molecule.
(iii) Continuous spectra - these are emitted from hot solids - the intensity /wavelength distribution
depending on the temperature of the solid.
Absorption spectrum
This type of spectrum is formed by the passage of light through a cooler vapour. The vapour then
absorbs those regions of the spectrum which it would have emitted had it been in an excited state.
In the cooler gas most of the atoms are in the ground state (the lowest possible energy state).
Absorption occurs when the electrons in the atoms absorb energy from the incoming radiation and
then reradiate it in all directions.
In 1814 Fraunhofer discovered that the spectrum of the Sun’s radiation was crossed by hundreds of
absorption lines which are called Fraunhofer lines. These are due to the absorption of light from the
centre of the Sun by the outer and cooler layers. Study of these lines gives the astronomers a way
of determining the composition of the Sun.
One of the simplest of spectra to explain is that of hydrogen and it is what we will consider next
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SPECTRA - THE KEY TO ATOMIC STRUCTURE
If we look at the spectrum of atomic hydrogen we
can see that it is made up of series of lines. This
arrangement of lines is unique to hydrogen, other
monatomic gases have a line spectrum but no other
element shows the same spectrum as hydrogen and
it is sensible to suppose that the spectrum somehow
reflects the atomic structure of the atom that
produced it.
Now in 1913 Neils Bohr suggested that the energy of
an electron in a hydrogen atom is quantised - that is
the energy can only have certain values. Each level is
given a quantum number.
When energy is given to the atom an electron can
jump up from one level to a higher level. This
transition will only take place if the energy provided is
equal to the difference in energy between the two
levels.
Energy level
Electron
transition
Energy may be supplied as either:
(a) radiation
(b) electrical
(c) a particle beam or
(d) heat
When an electron drops from one level to another a quantum of radiant energy is emitted and this
gives a line in the hydrogen spectrum. The energy of this quantum is given by the formula:
Quantum energy: E = hf where h is Planck’s constant and f is the frequency
The bigger the energy difference the greater the energy of the emitted quantum and therefore the
higher its frequency. Bohr actually derived a formula giving the energy of these levels and using it
we can calculate the values for the hydrogen atom shown in the diagram and the following table.
Bohr’s formula was: Energy of level n = 13.6/n2 eV.
n
1
2
3
4
5
6
Energy (eV)
-13.6
-3.4
-1.51
-0.85
-0.54
-0.38
Energy (J)
-21.8x10-19
-5.44x10-19
-2.42x10-19
-1.36x10-19
-0.87x10-19
-0.60x10-19
The energy level for n = 1 is called the ground state. Energy levels with n = 2 and above are called
excited states and the level with n = ∞ is known as the ionisation state.
Notice that all the energies are negative compared with the ionisation state which has zero energy.
This means that an electron in any of the other levels has to gain a certain amount of energy to
reach the ionisation level and so escape from the atom.
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