periodicity

• ionisation energy:

  • energy needed to remove the lost electron
  • when an atom loses an electron it becomes a positive ion (it had become ionised
  • evidence for existence of electron shells/energy levels
• first ionisation energy: the energy required to remove 1 mol of electrons from 1 mol of gaseous atoms.
• what is the general formula for a first ionisation energy reaction?

  $X(g) ->X^+$$(g)+$$e^-$
• second ionisation energy: the energy required to remove 1 mol of electrons from 1 mol of gaseous 1+ ions
• what is the general formula of second ionisation energy reaction?

  $X^+$$(g)->X^2(g)^+$$+$$e^-$
• as you descend G1, first ionisation decreases bc:

  • more shells/shielding increases
  • nuclear charge increases
  • more distance between electron + nucleus so electrostatic force is weaker
• as you go --> in the periodic table:

  • similar shielding, same energy level
  • nuclear attraction is greater bc of nuclear charge increasing
• ionisation energy is measured kJ mol-1 (kilojoules per mol)
• the lower the ionisation energy, the easier it is to form a positive ion
• what are the 3 factors that affect ionisation energy?

  1. Atomic radius 2. Nuclear charge 3. Shielding
• distance from nucleus: attraction rapidly falls off with distance
• nuclear charge: the more protons in the nucleus, the stronger the attraction of electrons
• shielding: inner shells repel outer shell electron, reducing the attraction between the outer shell electron and the nucleus
• 3 evidences for electron arrangement from ionisation energies:

  1. Down a group e.g. G2
  2. Across a period
  3. First ionisation energy (up to element 56)
• Down a group:

  • distance increases
  • more shielding
  • even though nuclear charge increases, distance + shielding means nuclear attraction decreases
  • less energy to remove the electron
• General increase:

  • similar shielding, however,
  • nuclear charge is bigger
  • nuclear attraction is stronger
  • more energy required to overcome

  G2-3 dip:

  • 3p1 on a new sub-shell
  • more shielding
  • nuclear charge increases, but shielding cancels it out
  • nuclear attraction decreases

  G5-6 dip:

  • nuclear charge increases but:
  • P- half full p sub-shell
  • S- 1p has both electrons, so it's filled up
  • due to electron repulsion, nuclear attraction is going to be weaker
  • less energy needed to overcome

  
• first ionisation energy - slower increase because d-subshell can hold 10 electrons in total. shielding is similar, nuclear attraction increases

  
• number of ionisation = number of electron removed
• across a period, b.p. increases because:

  • more strength of metallic bonding
  • larger no. of electrons
  • larger charge + smaller size of ions gives rise to a larger charge density
• the more electrons lost (e.g Al), the more strength in the bond
• Silicon:

  • high melting point as giant molecular structure like diamond
  • a lot of energy is required to break many covalent bonds
• P, S, Cl, Ar are simple covalent molecules. the larger the molecule, the greater the van der Waals' forces
• S8 - highest boiling point

  P4

  Cl2

  Ar - lowest boiling point

  • size increase
  • no. of electrons increase
  • london dispersion forces increase
  • more energy required
• atomic radius decreases across a period:

  • nuclear charge increases
  • similar shielding
  • nuclear attraction increases
  • greater attraction for electrons
  • pulls them in slightly
• positive ions (G1-3),

  ionic radius decreases,

  stronger nuclear charge,

  more electrons is removed,

  nuclear attraction stronger,

  making it smaller
• negative ions (G5-7)

  ionic radius increases

  more electrons

  nuclear attraction decreases

  increasing radius
• which group had the biggest ionic radius?

  group 5