Module 3.1.1- Periodicity

• Periodic trend in electronic configurations across periods 2 and 3

  Across period 2, 2s subshell fills first followed by the 2p subshell with 6 electrons

  Across period 3, the same pattern of filling is repeated for the 3s and 3p subshells
• Classification of elements into s-, p- and d-blocks

  
• First ionisation energy


  The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions
• Factors affecting ionisation energy

  Atomic radius

  Nuclear charge

  Electron shielding
• Effect of atomic radius on ionisation energy


  Larger atomic radius

  outer electrons further away from nucleus

  electrons less attracted to nucleus

  lower IE
• Effect of nuclear charge on ionisation energy


  More protons

  more positively charged nucleus

  stronger attraction between nucleus and electrons

  higher IE
• Effect of shielding on ionisation energy


  More electron shells

  more shielding

  less attraction between nucleus and electrons

  lower IE
• Successive ionisation energy of Fluorine


  The large increase between 7th and 8th ionisation energy suggests that the 8th electron must be removed from a different shell, closer to the nucleus and with less shielding

  
• Making predictions from successive ionisation energies

  allows to predict:

  number of electrons in outer shell

  group of element

  identity of an element
• Trend in first ionisation energy down a group

  First ionisation energies decrease down a group due to atomic radius increasing and more inner shells so shielding increases decreasing the nuclear attraction.

  
• Trend in first ionisation energy across a period


  First ionisation energy increases across a period.

  nuclear charge increases

  similar shielding

  nuclear attraction increases

  atomic radius decreases
• Which factor takes priority over another factor with the trend of ionisation energy across a period


  The increased nuclear charge is the most important factor for the general increase in first ionisation energy
• Why is there a fall in first ionisation energy from Beryllium to Boron

  It marks the start of the filling of the 2p subshell.

  2p subshell has a higher energy than the 2s subshell in Beryllium. Therefore the 2p electron is easier to remove than the 2s electrons.

  
• Why is there a fall in first ionisation energy from Nitrogen to Oxygen

  It marks the start of electron pairing in the p-orbitals of the 2p subshell. In oxygen, the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron. Therefore IE decreases from Nitrogen.

  
• Metallic bonding


  a bond formed by the attraction between positively charged metal ions and the electrons around them
• Structure of metallic bonding

  Giant metallic lattice which have delocalised electrons that are mobile with fixed cations in position.

  
• Properties of metals: electrical conductivity


  In solid and liquid states delocalised electrons are mobile, so they can move through the structure and carry the charge
• Properties of metals: melting and boiling points

  Melting point of metals depends on the strength of the metallic bonds. High temperatures are needed to overcome the strong electrostatic attraction between the cations and electrons, therefore have high melting and boiling points
• Properties of metals: solubility

  metals do not dissolve
• Giant covalent structures


  giant structures where atoms are held together by an array of strong covalent bonds
• Giant covalent lattice


  A three-dimensional structure of atoms, bonded together by strong covalent bonds.
• Example of giant covalent lattices


  Carbon, silicon

  both in group 4

  form 4 covalent bonds to its molecule

  tetrahedral structure 109.5 by electron-pair repulsion
• Properties of giant covalent structures: melting and boiling points

  high melting and boiling points

  strong covalent bonds so a large quantity of energy is needed to overcome the strong covalent bonds
• Properties of giant covalent structures: solubility

  insoluble in nearly all solvents

  covalent bonds holding atoms are too strong to be broken by interaction with solvents
• Properties of giant covalent structures: electrical conductivity

  non-conductors of electricity except graphene and graphite since there are delocalised electrons between layers (spare electrons come from the 3 bonds formed only by carbon instead of 4)
• Trend in melting points across periods 2 and 3

  Sharp decrease in melting point marks a change from giant to simple molecular structures

  
• Trend in structure across periods 2 and 3


  GMS- Li, Be, Na, Mg, AlGCS- B, C, SiSMS- N2, O2, F2, Ne, P4, S8, Cl2, Ar