Chapter 7

• What is the difference between the old (mendeleev) and new periodic table?

  Old - was arranged in order of increasing atomic mass

  New - elements are arranged in increasing atomic (proton) number
• Define periodicity?


  A repeating trend in properties of elements across each period:

  electron configuration

  ionisation energy

  structure

  melting points
• Define nuclear attraction?


  The electrostatic attraction between the positively charged nucleus and negatively charged electrons within the atom.
• Define ionisation energy?


  The energy required to remove one electron from each atom in one mole of gaseous atoms of an element.
• Factors affecting ionisation energy

  Atomic radius - greater distance between nucleus and electrons = weaker attraction

  Nuclear charge - more protons = stronger attraction

  Inner shell shielding - inner electrons repel outer ones, more shells = weaker attraction
• First ionisation energy trends in the periodic table

  Decrease going down a group - increases atomic radius, more inner shells, weaker attraction

  Increase across a period - nuclear charge increases, shielding remains constant, slight atomic radius decrease , stronger attraction
• Ionisation energies in atomic orbitals

  2p subshell has a higher enegry than the 2s subshell

  Therefore it is easier to remove the outer 2p electrons in Boron than the 2s electron in Be - further from the nucleus = weaker attraction

  Spin paried electrons repel each other - easier to remove paired electrons
• Successive ionisation energies


  Equations MUST include state symbols (always GASEOUS)

  removing all electrons from an atom

  outer shell electrons are removed one at a time

  as electrons are removed attraction increases because atom is smaller

  He(g) —> He+(g) + e- first ionisation energy

  He+(g) —> He2+(g) + e- second ionisation energy
• Define metallic bonding


  The strong electrostatic attraction between metal cations and delocalised electrons
• Metallic lattice structure


  Electrons can move freely and form a "sea" of delocalised electrons

  Positively charged metal ions in a regular arrangement
• Properties of metals


  High boiling point - metallic bonds are strong, relying on charge of cation and number of delocalised electrons

  (exeption = mercury)

  High electrical conductivity (solid and in molten form) - sea of delocalised electrons
• Properties of non-metals


  Structures:

  Giant covalent (boron, silicon and carbon - allotropes) - regular arrangement of atoms all covalently bonded

  Simple molecular lattices - strong covalent bonds, weak London forces, low melting and boiling points, cannot conduct electricity, like dissolve in like substances (polar vs non-polar)
• Giant covalent structures


  Carbon (diamond form) and silicon use 4 outer shell electrons to form covalent bonds to other carbon/silicon atoms

  Tetrahedral structure - 109.5 bond angle

  very high melting / boiling point

  insoluble

  does NOT conduct electricity
• Graphite


  trigonal planar - 120 bond angle

  3/4 outer electrons bonded

  delocalised electrons (sea)

  form hexagonal layers

  very high melting and boiling points

  insolbule

  does conduct electricity
• Graphene


  single layer of graphite

  can still conduct electricity

  very strong

  very thin
• Periodic trends in melting points

  giant structures have stroner forces to overcome - higher melting points

  simple molecular structure have weaker forces to overcome - lower melting points