P block

• An atom must exist for >10^ -14 seconds to be classed as an element

• Metalloids

  • Have properties of both metals and non-metals
  • Can form cations of the atoms
• Ionisation energy


  • The energy required to remove completely an electron from the gaseous atom or molecule in its ground state
  • Values are always positive
  • Units are kJ mol-1
• Across the periodic table - IE increases

  • Number of protons increases
  • Effective nuclear charge increases
  • Harder to separate an electron from the nucleus
• Down the periodic table - IE decreases

  • Electrons are in valence shells which are further away from the nucleus
  • Less of a stabilising effect
  • Easier to remove an electron
• The loss of an electron is slightly favourable for oxygen due to electrostatic repulsions between electrons
• Group 13 elements Gallium and Thallium have higher IE values than expected. Gallium is preceded by the first set of d orbitals and Thallium is preceded by the first set of f orbitals
• d and f electrons are weakly shielding, resulting in a higher effective nuclear charge than expected
• Relativistic effects impact the ionisation energy of Thallium
• Electron affinity


  The energy released when a gaseous atom, molecule or ion in its ground state gains an electron
• Electron affinity

  • Values mostly positive, therefore negative enthalpy
  • Measured in kJ mol-1
• Fluorine + Oxygen EA

  Very small positive electron affinity, as the electron-dense system experiences electrostatic repulsions that cancel out some of the positive stabilising effect
• Group 13 EA

  D block contraction - Gallium is slightly smaller than it would otherwise be, the added electron feels a higher effective nuclear charge and is more stable

  General increase in electron affinity with a slight dip for group 15
• Electronegativity (χ)


  The ability of an atom to attract electron density towards itself in a molecule
• Electronegativity

  • Pauling's scale: 0 ≤ χ ≤ 4
  • Increases across the periodic table up to halogens
  • Decreases down the periodic table due to more shielding
  • May vary by hybridisation, with electrons in orbitals with more s character being held more tightly (sp > sp2 > sp3)
• Van Arkel Ketelaar triangle

  • Ionicity parameter - the difference in electronegativity of the 2 atoms in the binary compound
  • Covalency parameter - the average electronegativity of the 2 atoms in the binary compound

  
• Slater's rules

  Can be used to estimate the Zeff (effective nuclear charge)

  Don't count the electron in consideration or electrons in higher shells

  • s/p electrons: Electrons in with the same principle quantum no. - +0.35, Electrons in the (n-1) shell - +0.85, Electrons in the (n-2) shell - +1.0
  • d/f electrons: Electrons with the same principle quantum no. - +0.35, Anything in a lower shell contributes 1.0
• Effective nuclear charge

  Increases across the period
• Effective nuclear charge

  Increases down the period
• Covalent radius


  Half the length of a symmetrical homonuclear bond
• Metallic radius


  The equivalent distance between ions in a metal lattice
• Ionic radii


  Increases with -ve charge and decreases with +ve charge
• Inert pair effect


  The tendency of electrons in the outermost atomic s orbital to remain unionised/unshared in groups 13-16
• If a bond's orbital overlap is good

  The energy to promote an electron will be low
• Formation of the bond is exothermic

  Offsets the endothermic 'cost' of sp3 hybridisation
• Relativistic effect


  Theory of relativity - fast things weigh more

  In heavier atoms, electrons can move very fast and increase in mass

  Causes the orbitals to contract - lower energy

  Most contraction in s orbitals > p > d + f
• Direct relativistic orbital contraction

  An electron can move so fast its mass contracts the orbital

  The orbital energy is lower

  Larger effective nuclear charge
• Bonding

  • Orbitals become more diffuse going down the group - longer bonds - weaker
  • First row elements form multiple bond (double/triple bonds)
  • Small atoms (short bond) + good orbital overlap
  • Able to have pi interactions
• P𝛑-d𝛑 bonds

  A filled p orbital can donate electron density to another atom's empty d orbital
• Favoured O.S.

  $(n - 2)$

  n = group no.

  
• Wade's rules - Boranes + carboranes (CH+)

  • Each B-H (C-H) contributes an electron pair
  • Each spare H contributes 1 electron
  • +/- any charge on the cluster
  • PSEPs = no. electrons/2
  • No. vertices = PSEPs - 1
  • if no. vertices = no. BH units - Closo
  • if no. vertices = no. BH units - 1 - Nido
  • if no. vertices = no. BH units - 2 - Arachno
• Silicone

  • Small covalent radii + good orbital overlap - Pi bonding
  • Terminal monomers have 3 R groups
  • Monomers with 3 hydroxyl groups are able to cross link chains
  • p 𝛑-d 𝛑 bonding

  
• Pauling's rules - reactivity of oxyacids


  Acid's pKa = 8 - (5 x E=O bonds)
• Group 13 - the Triels

  Chemistry is dominated by electron deficiency

  Act as Lewis Acids

  Full p orbital on hallide donates electrons density into vacant p orbital on Boron

  Dimerisation is preferred to 𝛑 bonding for larger group 13 elements (Al2Cl6)
• Boranes (B2H6)

  Borons are tetrahedral (4 x sp3 orbitals)

  • 8 electrons are used for the B-H bonds (2 centre-2 electron bonds)
  • 4 electrons are used for the B-H-B bonds (3-centre-2 electron bonds)

  
• Group 14 - the Tetrels

  Hydrides become more stable going down the group

  • larger, more diffuse orbitals
  • more polar bonds
  • metal is less shielded
  • low-lying LUMOs