Module 2. Electrons, bonding and structure. 1 - 4

Ionisation energy
The energy required to form postive ions is known as ionisation energy. Ionisation energies provied evidence for a model of the atom in which electrons are arranged shells.
The first ionisation energy of an elemeent is the energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseuos 1+ ions.

Factors affecting ionisation energy

Electrons in the outer shell are removed first since they experience least nuclear attraction. The outer-shell are furthest away from the nucleus require the least ionisation energy. The nuclea attraction experienced by an election depends on 3 factors:

• Atomic radius: The greater the atomic radii, the smaller the nuclear attraction experienced by the outer electrons
• Nuclear charge: The greater the nuclear charge, the greater the attractive force on the outer electrons
• Electron shielding: Inner shells of electrons repel the outer-shell electrons. This repelling effect is called electron shielding. The more inner shells there are, the larger the shielding effect and the smaller the nuclear attraction experienced by the outer electrons. 

Successive ionisation energies
Successive ionisation energies are measure of the energy required to remove each electron in turn. For example, the second ionisation energy is a measure of how easily 1+ ions loses an electron to form a 2+ ion.
An element has as many ionisation energies as it has electrons. For example, Lithium has 3 electrons therefore it has 3 successive ionisation energies.

Each successive ionisation energy is larger than the one before.

• As each electron is removed, there is less repulsion between the electrons and each shell will be drawn in to be slightly closer to the nucleus.
• As the distance of each electron from the nucleus decreases slightly, the nuclear attraction increases. More ionisation energy is needed to remove each successive electron.  

Evidence for shells
Sometimes you will see a large increase in successive ionisation energies. An example is between the 1st and 2nd ionisation energies Lithium. This large increase shows that the second electron has been removed from a different shell, closer to the nucleus and with less shielding from inner electrons. 

A shell is a group of atomic orbitals with the same principal quantum number, n. Also known as the main energy level.

Atomic orbitals
An atomic orbital is a region within an atom that can hold up too 2 electrons with opposite spins. There are 4 different types of orbital - s, p, d and f. Each has a different shape.

S-orbitals
An s-orbital has a spherical shape. From n=1 upwards, each shell contains one s-orbital. 2 s electrons in each shell.

P-orbitals
A p-orbital has a 3-dimensional dumb-bell shape. From n=2 upwards, each shell contains 3 p-orbitals, px, py and pz, at right angles to one another. This gives a total of 6 p electrons.

D-orbitals and f-orbitals
The structure of d- and f-orbitals are more complex. From n=3 upwards, each shell contains 5 d-orbitals, this gives a total of 10 d electrons. From n=4 upwards, each shell contains 7 f-orbitals. This gives a total of 14 f electrons. 

Electrons within an orbital always have oposite spins. This is called Hund's rule.

Sub-shells

A sub-shell is a group of the same type of atomic orbitals (s, p, d and f) within a shell.

Electron energy levels
The sub-shells within a shell have different energy levels. The order of these energy levels is shown in the figure below. Within a shell, the sub-shells energies increase in the order of s, p, d and f.
An electron configuration is the arrangement of electrons in an atom. 

Electron shells overlap

• The larger the value of the principal quantum number n, the higher the energy level and the further the shell is away from the nucleus
• As we move from one shell to the next, a new type of sub-shell is added.
• Within a shell, the sub-shell energies increase in the order s, p, d and f. 

Need to take care when you go beyond the 3p sub-shell. Orbitals fill in the energy level order, but:

• the 4s energy level is below the 3d energy level.
• the 4s-orbitals fill before the 3d-orbitals 

Therefore, the 4th shell starts to fill before the 3rd shell has completely filled. For example the electron configuration of Potassium would be 1s2, 2s2, 3s2, 3p6, 4s1. 

Remember, 4s: first in, first out.

Sub-shells and the periodic table
The periodic table is structured in blocks of 2, 6, 10 and 14 elements, linked to sub-shells. The pattern mirrors the sub-shells that are being filled. We can easily see the pattern by dividing the Periodic Table into blocks.

Shortening an electron configuration
Long electron configuration can be abbreviated by basing the inner shell configuration on the noble gas that comes before the element in the Periodic Table. The shortening also allows you to concentrate on the important outer shell portion of the electron configuration. It is these electrons that are responsible for the chemical character of the element.

Electron configuration of ions
When positive ions are formed, electrons are removed from the highest energy orbitals.
When negative ions are formed, electrons are added to the highest energy orbitals. One exception to his rule is the elements Sc to Zn at the top of the d-block. You will remember that the 4s sub-shell is filled before the 3d sub-shell.
The 4s and 3d energy levels are so close together that, after the 4s-orbital has been filled, it is actually at a higher energy than the 3d level. So, 4s-electrons are lost before the 3d level.