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.
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.
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.