Energy levels
Electrons are located in fixed energy levels. All atoms have their own unique set of energy levels which depend on how many protons and neutrons an atom contains.
Energy levels are numbered from 1 up to infinity. The lowest energy level, 1, is located closest to the nucleus. The further out the energy level is, the less attracted it is to the nucleus.
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Orbitals and sub-levels
Electrons occupy areas called orbitals. However, it is not possible to locate exactly where an electron is: an orbital is simply the space in which there is a high chance an electron will be.
Orbitals differ in shape:
s-orbitals: these are spherical and found in all energy levels. An s-orbital is named after which energy level it is located in. For example, one located in the first energy level is called a 1s orbital.
p-orbitals: these have a 3D figure of eight shape and exist as groups of three. They are found in their groups of three in every energy level apart from the first. Each p-orbital within the same energy level has the same energy but a different orientation (x, y and z). A p-orbital is named in the same way as an s-orbital. So, a p-orbital located in the second energy level is called 2p (2px, 2py, 2pz).
The third and subsequent energy levels contain five d-orbitals. Seven f-orbitals are then found in the forth and subsequent energy levels and so on. These orbitals each have their own shape.
S, p and d orbitals have different energies. No matter the energy level, s-orbitals have the lowest energy, p-orbitals a higher energy, and then the energy continues to increase from d-orbitals onwards. Therefore, each energy level is split of into sub-levels of different energies.
A summary of the orbitals in each energy level can be found in the table below.
|
Energy level |
Number and type of orbital |
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|
1st sub-level |
2nd sub-level |
3rd sub-level |
4th sub-level |
5th sub-level |
|
|
1 |
1 x 1s |
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|
2 |
1 x 2s |
3 x 2p |
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|
3 |
1 x 3s |
3 x 3p |
5 x 3d |
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|
4 |
1 x 4s |
3 x 4p |
5 x 4d |
7 x 4f |
|
|
5 |
1 x 5s |
3 x 5p |
5 x 5d |
7 x 5f |
9 x 5g |
Shells
Due to the fact that sub-levels carry different energies and that as the energy level number increases the levels become closer together, high energy and low energy sub-levels inevitably overlap.
Many sub-levels have similar energies and are grouped together to form a shell.
The table below illustrates the arrangement of shells and the largest number of electrons possible in each shell.
| Shell number | Orbitals in shell |
| 12345 | 1 x 1s1 x 2s, 3 x 2p1 x 3s, 3 x 3p1 x 4s, 5 x 3d, 3 x 4p1 x 5s, 5 x 4d, 3 x 5p |
Electrons
Electrons repel each other therefore it is not possible to have more than two in one orbital.
Two electrons spinning within the same orbital in opposite directions creates a magnetic field and small magnetic attraction between the electrons.
It is, therefore, possible to calculate the maximum number of electrons possible within a sub-level, as the table below illustrates.
| Shell | Number of electrons in each sub-level | Maximum number of electrons |
| 12345 | 2 x 1s2 x 2s, 6 x 2p2 x 3s, 6 x 3p2 x 4s, 10 x 3d, 6 x 4p2 x 5s, 10 x 4d, 6 x 5p | 2881818 |
Electron arrangement
How electrons fill orbitals is dictated by three rules:
- The Aufbau/building principle: the lowest energy orbitals are always filled first.
- Hund’s rule: until all orbitals of the same energy are singly occupied, electrons never pair up within the same orbital and all electrons which are unpaired spin in a parallel fashion.
- Pauli exclusion principle: only a maximum of two electrons can occupy the same orbital and they spin in opposite directions.
The name given to the way in which electrons are arranged within an atom is called electronic configuration. This arrangement can be illustrated using two methods:
With the arrow and box method the orbital is represented as a box and the electrons as arrows. The direction of the arrow represents the direction of spin.
The orbital method represents the number of electrons in an orbital in superscript.
In shorthand (for both methods) full shells are represented by the element’s symbol in a square bracket.
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Electron arrangement in ions
In order to figure out the electronic configuration of ions you need to add or remove the correct number of electrons. This number can be worked out by using the following rules:
- first remove the outer shell electrons
- remove the electrons in this order: p-electrons then s-electrons then d-electrons
- paired electrons must be removed before unpaired electrons within the same sub-level
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Effect of electronic configuration on chemical properties
The attraction between the nucleus and the outer electrons dictates the chemical properties of an atom. This attraction depends on both the electronic configuration and the proton number.
The neutron number does not have an effect on an atom’s chemical properties as neutrons do not have a charge and so no attractive force on the nucleus. Therefore, isotopes, which differ in neutrons and not in protons, have similar chemical properties.