Atomic structure

Ionisation is the removal of one or more electrons.

The first ionisation energy is the energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions.

X(g) = X+ (g) + e-

Ionising atoms or molecules is an endothermic process - energy needed.

The lower an ionisation energy, the easier it is to form an ion.

Nuclear charge affects ionisation energy because the more protons in the nucleus, the greater te +ve charge in the nucleus and the stronger the attraction is for the electrons.

The distance of the outermost electron from the nucleus also affects the ionisation energy because attraction between E and nucleus falls rapidly with distance.

Shielding also affects ionisation energy as the number of electrons between the outer electrons and the nucleus increases, the outer electrons feel less attraction to the nuclear charge. This lessening of the pull of the nucleus by the inner shells of electrons is known as shielding.

Successive ionisation energies involve removing additional electrons.

Each time you remove an electron, theres a successive ionisation energy.

The second ionisation energy is the energy needed to remove 1 electron from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions.

X(n-1) (g) = X(n+) (g) + e-

Successive ionisation energies show shell structure.

Within each shell, successive ionisation energies increase. Electrons are being removed from an increasingly +ve ion - less repulsion amongst the remaining electrons, they're held more strongly by the nucleus.

The big jumps in ionisation energy happen when a new shell is broken into - an electron is being removed from a new shell closer to the nucleus.

Graphs can tell you which group of the periodic table an element belongs to - just count how many electrons are reomved before the first big jump.

Graphs can also help predict electronic structure - working from left to right, count how many points there are before each big jump to find out how many electrons are in each shell, starting with the first.

• The first ionisation energies of elements down a group of the periodic table decreases.
• The first ionisation energies of elements across a period generally increase.

IONISATION ENERGY DECREASES DOWN GROUP 2

• If each element down group 2 has an extra electron shell compared to the one above, the extra inner shells will shield the outer electrons from the attraction of the nucleus.
• The extra shell means that the outer electrons are further away from the nucleus - the nucleus's attraction will  be greatly reduced.
• These factors will make it easier to remove outer electrons - resulting in a lower ionisation energy.

IONISATION ENERGIES INCREASE ACROSS A PERIOD

• As the number of protons is increasing there is a stronger nuclear attraction.
• All the electrons are at roughly the same energy level. This means there is generally little extra shielding effect or extra distance to lessen the attraction from the nucleus.

There are small drops between groups 2&3 and 5&6.

THE DROP BETWEEN GROUPS 2 & 3 SHOWS SUB SHELL STRUCTURE

• Al's outer electron is in a 3P orbital rather than a 3S. 3P has a slightly higher energy, so the elctron is, on average further away from the nucleus.
• The 3P orbital has extra shielding provided by the 3S2 electrons.
• Both of these factors are strong enough to overcome the effect of the increased nuclear charge - resulting in a drop in ionisation energy.
• This provides evidence for the theory of electron sub-shells.

THE DROP BETWEEN GROUPS 5 & 6 IS DUE TO ELECTRON REPULSION

• The shielding is identical in Phosphorus and Suflur atoms, and the elctron being removed is from an  identical orbital (3P)
• In the case of Phosphorus, the electron in the 3P orbital is singular. In Sulfur, the  orbital has 2 electrons in it - its fully occupied.
• The repulsion between the pair of electrons means that it is easier to remove from shared orbitals.
• This provides more evidence for the electronic structure model.