Periodicity

Periodicity

Atomic radius
1. Decreases in size the further to the right the element is along a period due to there being a larger positive charge in the atoms nuclei
2. Increases in size the further down a group an element is, as there are more electronic shells in elements further down the group
Ionic radius
1. Anions (Ions with a negative charge) have a larger radii than their atoms counterparts due to them having a larger negative charge, making the attraction to the nuclei weaker
2. Cations (Ions with a positive charge) have a smaller radii than their atom counterparts due to them having a larger positive charge, making the attraction to the nuclei stronger
Ionisation energy
1. Ionisation energy increases the further the element is to the right in a period because the nuclear charge increases, attracting the negative electrons more to the positive nuclei, The atomic radii decreases, therefore the electrons are closer and more electrons are added to the same shell.
2. Ionisation decreases the further an element is down a group because; The atomic radii decreases, therefore the electrons are further away; The nuclear charge increases but the electrons are further away so there is a weaker attraction; and the electrons in lower shells block attraction causing electron shielding
3. Anomalies to this rule:
  1. 1. Boron has a lower IE than Beryllium because the 2p electrons are slightly higher in energy than the 2s electrons and so the ionisation for Boron is lower than for Beryllium (Moving from an S shell to a P shell lowers the IE)
  2. 2. Oxygen has a lower IE than Nitrogen because the Px, Py and Pz only contain one electron. The extra electron in Oxygen causes a pair and repulsion making it easier to remove, hence giving it a lower IE than Nitrogen (When the P-shell has one electron in all 3 orbitals, the next electron [4p^4] will have a lower IE due to that fourth electron being easier to remove)
Electron Affinity
1. The energy released when 1 mole of electrons is attached to one mole of neutral atoms or molecules in the gas phase
2. X(g) + e- ===> X-(g)
Electronegativity
1. In general, the electronegativity increases the further to the right in a period the element is, and decreases the further down a group it is.
2. Electronegativity is the tendency of an atom to attract a shared pair of electrons towards itself. An atom's electronegativity is affected by both its atomic number and the distance at which its valence electrons reside from the charged nucleus.

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