Week 6: Periodicity and S-block

Periodicity

Trends in chemical and physical properties across periods or down groups

Atomic radius

Half the distance between two nuclei

Ionic radius

The internuclear distance between ionic solids

Ionisation energy

The energy change when an electron is removed from an atom of the element in the gaseous phase

Electron gain energy

The energy change that occurs when an electron is attached to an atom in the gaseous phase

Electronegativity

The tendency of an atom to attract shared electrons towards itself within a chemical bond

Trends across the periods: Atomic/Ionic Radius

• Decreases from left to right

• Higher proton number while general atom size remains similar

• Z_eff higher, valence electrons contract, shorter distance result

• Ne is much higher due to its gaseous state, it is measured in Van Der Waals

Trends going up the groups; Atomic/Ionic Radius

• Increase from top to bottom

• Increasing period means increase in principle quantum number of valence orbital

• Valence orbital becomes larger

• Extra orbital are much further away from the nucleus

Ionisation energy across the period

• ionisation energy increases

• Due to valence electron closer to the nucleus as you move from left to right

• Hence there are more electrostatic attraction requiring more energy to overcome so higher ionisation energy

Ionisation energy up the group

• Ionisation energy decreases

• Due to valence electron being further orbit

• Weaker attraction between valence electron and nucleus

• More shielding from orbitals

• Less energy to overcome so lower ionisation energy

Why does the reaction between group 1 and water becomes more vigorous down the group

• Principle quantum number of valence shell increases

• Valence electron further away from nucleus

• Easier to remove an electron

Group 2

Group 1 and 2 in society