The Bohr atom and the quantum mechanical model of the atom
The Bohr atom
The model of the atom we are used to dealing with in GCSE chemistry is called the Bohr atom;
named after Neils Bohr
a Nobel Prize winning scientist. This model of the atom assumes a
small nucleus with the electrons
orbiting the nucleus in circular orbits. Bohr placed the
electrons in energy levels or
shells. This picture of the atom
should be familar to you from your gcse science course. The diagram opposite is probably very
familiar to you,
it shows a nucleus which contains the protons and neutrons with
the electrons in their shells
orbiting the nucleus in circular orbits. The
first shell holds 2 electron and
the remainder of the shells holding 8 electrons
.
Bohr calculated the energies of these electron orbits or
shells and he suggested that
there were only certain allowed energy levels. However
Bohr's model does not offer a completely accurate description of how the electrons
are arranged
in atoms. In fact this model of the atom can only explain some of the observations
produced from emission spectra for atoms or ions which contain a
single electron. A new model to help explain some of the observations for multi-electron atoms
was put forward by brilliant physcist Edwin Schrodinger in 1926 based on what we now call the quantum
mechanical model of the atom.
Quantum numbers
The wave equations produced from the Schrodinger equation produce 3 quantum numbers that offer a description of the location,
energy and orientation in space of the electrons inside the atom. One of these
quantum numbers, the principal quantum number is similar to
Bohr's principal energy levels.
- The principal quantum number (n)- It is perhaps
easy to think of the principal quantum number as the shell
or energy level for Bohr atom model we used in gcse chemistry.
Exactly as in gcse science the principal quantum number (n) can have integral values of 1,2,3,4,5.....etc. As n increases
the energy levels or shells are getting larger and the electron
is found further from the nucleus.
This means that the electrons are increasing in energy.
- The other quantum numbers provide information on the location,
energy and orientation in space of the electrons inside the atom.
However when the emission spectra of the
elements are studied closely it is found that the main principle energy levels
within atoms are in fact split into smaller sub-shells or sub-levels.
Evidence from the ionisation energies of the various
elements
also help confirm the existence of these sub-shells. So a better model for the electronic structure of
atoms might be
similar to the one below, this shows an atom of chlorine.
sub-shells and sub-levels
Quantum theory as well as
evidence from ionisation energies and emission spectra
suggests that the main principal energy levels are in
fact split up into sub-levels or sub-shells with slightly different energies. These sub-shells
or sub-levels are described by the letters: s, p, d, f and g. These letters are
taken from names used to describe parts of the emission spectra of
atoms, s for sharp, p for principal and d for diffuse . Each
sub-shell contains a number
of orbitals. Orbitals
are a region of space where there is a high probability (95% chance) of finding the
electron. Each orbital
can hold a maximum of 2 electrons with opposite spins. The number of
orbitals increases
as the principal quantum number increases, this means that the number of electrons
that each shell can hold increases as the shells
become larger. We can use the formula 2n2 to calculate the number of electrons
present in each shell (n is simply the
shell number).
This is summarised in the table below:
| |
First shell/energy level |
Second shell/energy level |
Third shell/energy level |
Fourth shell/energy level |
| sub-shells present |
s |
s, p |
s, p, d |
s, p, d, f |
| Number of orbitals present |
1 |
1, 3 |
1, 3, 5 |
1, 3, 5, 7 |
| Maximum number of electrons present |
2 |
2, 6 |
2, 6, 10 |
2, 6, 10, 14 |
As the value of n increases the energy of the shells or
principal energy level increases as the electrons get ever further
from the nucleus. However the gap between different
shell decreases as we add more shells, for example there is a larger
energy gap between shells 1 and 2 then there is between shells 3 and 4. As a consequence of this
the the sub-levels
in different shells start to overlap as shown in the diagram above. You may notice for example that
the 3d orbitals have more energy
than the 4s orbital despite being in the third
shell.
You will not doubt have also noticed in the diagram opposite that the energy level diagram for a
single electron atom such
as hydrogen is different from that for multi-electron atoms.
In hydrogen all the sub-levels within any particular shell
are all degenerate, clearly from the diagram for
multi-electron atoms this is not the case and the sub-levels in any
particular shell have different energies.
Key Points
- The Bohr model of the atom assumes the
electrons are particles that orbit the
nucleus in fixed energy shells or rings
- The quantum mechanical model of the atom assumes the
electron is a wave and not a particle.
- Like all waves its motion and properties can be described by a series of wave equations. These
equations can be used to describe the
energy, motion and behaviour of the electrons in an
atom.
- The quantum mechanical model of the atom has the
shells split up into smaller sub-shells or sub-levels.
- The electrons in the new model of the
atom do not orbit the nucleus in circular orbits as described by Bohr. Instead
they are found inside orbitals. Orbitals
are 3-d shapes centred on the nucleus where there is a
high probability of finding the electron.
- Each orbital can hold a maximum of 2 electrons with opposite spins.
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