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Applied Chemistry Fundamental
Physical Chemistry

Electron Arrangement

Electron Arrangement

This topic often confuses people so don't worry if you don't understand it first time. Read up on electron arrangment in textbooks and get other people to explain it to you.

Before this level of study you will have been told that electrons have fields of 2,8,8 maximum capacity. Well, this is not true. As well as main energy levels, there are also sub-shells. The table below outlines this information.

Main LevelSub-levelElectron capacity

For these sub-shells it is important to note that the 4s requires less energy for electrons to enter it than the 3d. Other than this the electrons fill up each successive sub shell in order to form electrons. So below are the electron configurations of some elements.

some electron configurations

On the image, highlighted in blue are two features:

  • The first is that at Scantium (Sc) the electrons in 3d start to fill up after 4s.
  • The second is [Ne], this is a shorthand way of writing the electronic configuration, the bit highlighted means the electron configuration of Neon then the rest. You can use any noble gas dependent on its size.

    Hund's Maximum Multiplicity Rule

    This rule is an important one when considering electron arrangement and ionisation energy. It states that in subshells, the electrons will as far as possible be unpaired before; this is to be as energy efficient as possible. It is demonstrated below using spin diagrams.

    hunds rule of maximum multiplicity

    Ionisation Energy

    The first ionisation energy is defined as: The energy/enthalpy change when one mole of electrons is removed from one mole of a gaseous element, the equations for the first and second ionisation energies are shown below.

    equations for the first and second ionisation energies

    The ionisation energies of elements have a distinctive pattern, which provide evidence for the structure of electrons that is explained above.

    Group II patterns

    first ionisation energies of group 2 elements

    Down the groups there is a decrease in first ionisation energy. This is because of the following factors:

    Atomic Radius. This is the distance of the outermost electron to the nucleus. As you go down the groups, this distance gets bigger, so it becomes easier to remove electrons (less energy is required) as they are further from the attractive force of the nucleus.

    Electron Shielding. As you go down the group there is a new n (main) electron level. This means that other electrons are below the outermost. Because of this, there is a degree of repulsion, which means the electron is more readily lost.

    Period III Patterns

    first ionisation energies of period 3 elements

    The pattern of first ionisation energies across a period is a little more complex, because it does not show a perfect pattern. However, there is a general increase along the period because of increasing nuclear charge the force of attraction from the increasing number of protons means the electrons are being held in place with more force so more energy is required to remove them.

    There is a drop in first ionisation energy from magnesium to aluminium, this is because from Al, the electrons start filling the 3p sub-shell, and because this is in a higher energy level, it is lower.

    The next slight fall is from Phosphorus to Sulphur. This is because of Hund's maximum multiplicity as explained above. The electrons begin to pair up, and so less energy is required to remove it since the two electrons in a pair repel each other slightly.