Week 3

• Properties of metals:

  • Conduct heat & electricity well
  • Shiny
  • Lose valence electrons to form cations
• What happens when a transition metal compound forms?


  Relative energies of 3d & 4s energy levels swap
• How do u calculate d block electron count?


  Group number - metal charge
• P-block metal compounds act as if their d electrons are core (d electrons lower in energy than 4s & 4p)
• Properties of non-metals:


  Hard to remove valence electrons

  Gain electrons to form anions
• Properties of non-metals near metal/non-metal boundary:

  • Hold valence electrons strongly (metallic behaviour supressed)
  • Not good at attracting electrons or making cations/anions
  • Share electrons with other species
• What is electronegativity?


  Ability of atom to attract electrons towards itself when interacting with another atom.
• Electronegativity trend:


  Increases across periodic table

  Decreases down group
• What is the chemical behaviour of carbon?


  Hard to ionise
• What is the chemical behaviour of silicon?


  Non-metallic & reluctant to gain/lose electrons
• What is the chemical behaviour of germanium?


  Metallic & forms Ge (4+) with difficulty
• What is the chemical behaviour of tin?


  Metallic & forms Sn (4+) but Sn (2+) also possible
• What is the chemical behaviour of lead?


  Metallic & mainly forms pb (2+)
• Properties of unreactive elements:

  • Completely filled valence energy level
  • hard to remove/share electrons as too tightly held
  • Effective nuclear charge very high
  • Proportion of core (shielding) electrons minimised
  • Hard to gain electrons as they would have to enter next energy level (further from nucleus)
• What does the shape of an orbital tell us?


  Where electron most likely to be found
• What does shading of the orbital tell us?


  Sign (+ or -) of the wavefunction used to generate the shape

  A.K.A (the phase of the wavefunction)
• When do we need to know the sign of the wavefunction?


  When bringing together orbitals from different atoms to make bonds
• What is constructive overlap in the context of orbitals?

  2 atomic orbitals bond covalently to make a molecular orbital

  (electron density between atoms increases)
• What is destructive overlap in the context of orbitals?


  2 atomic orbitals create an anti-bonding orbital by shoving eachother apart

  (no electrons halfway between the atoms)
• Overlap rules:

  • Only orbitals with similar energies can overlap (valence)
  • For every bonding orbital created, a higher energy anti-bonding orbital also created
  • Orbital 'symmetries' must match for overlap to occur
• Why is ammonia stable?


  3 bonding pairs of electrons & 1 lone pair that isn't bonding or anti-bonding
• Why might a molecule fall apart?


  Anti-bonding energy level fills with electrons
• What is effective nuclear charge (Z*)?


  Charge experienced by outer electrons when shielding taken into account.

  As % of core electrons decreases, Z* increases