Module 3

• s, p, d blocks


  Classification of elements according to which orbitals the highest energy electrons are in
• Elements


  • Classified as s, p or d block
  • Highest energy electrons are in different orbitals
• Atomic radius decreases from left to right across a period
• Exactly the same trend in period 2
• Periodicity


  A repeating pattern across different periods
• Properties that display periodicity

  • Atomic radius
  • Melting points
  • Boiling points
  • Ionisation energy
• Elements are arranged in increasing atomic number in the periodic table

  
• Elements in Groups have similar physical and chemical properties
• Elements in periods show repeating trends in physical and chemical properties
• Period 2


  • Li
  • Be
  • B
  • C
  • N
  • O
  • F
  • Ne
• Period 3


  • Na
  • Mg
  • Al
  • Si
  • S
  • Cl
  • Ar
• First ionisation energy


  Energy needed to remove an electron from each atom in one mole of gaseous atoms
• The equation for 1st ionisation energy always follows the same pattern
• Factors affecting ionisation energy

  • Attraction of the nucleus
  • Distance of electrons from nucleus
  • Shielding of nuclear attraction
• Successive ionisation energies give information about electronic structure
• Helium has the largest first ionisation energy
• Metallic bonding


  The electrostatic force of attraction between the positive metal ions and the delocalised electrons
• Factors affecting strength of metallic bonding

  • Number of protons/Strength of nuclear attraction
  • Number of delocalised electrons per atom
  • Size of ion
• Types of bonding and structures

  • Covalent: shared pair of electrons, macromolecular
  • Metallic: electrostatic force between positive ions and delocalised electrons, giant metallic lattice
• Metals are malleable as the positive ions in the lattice are all identical, so the planes of ions can slide easily over one another
• Metals are conductive when solid due to the delocalised electrons that can move through the structure
• Metals are conductive when molten
• Metals are shiny solids
• Trends in melting and boiling points

  • Na, Mg, Al - high due to strong metallic bonding
  • Si - very high due to strong covalent bonds in macromolecular structure
  • Cl2, S8, P4 - low due to weak intermolecular forces between simple molecules
  • Ar - low as monoatomic with only weak London forces
• Similar trends in period 2: Li, Be - metallic bonding (high mp), B, C - macromolecular (very high mp), N2, O2 - molecular (gases, low mp), Ne - monoatomic gas (very low mp)
• Atomic radius


  Increases down the Group
• Atomic radius


  • As one goes down the group the atoms have more shells of electrons making the atom bigger
• Electronic Structure


  Group 2 metals all have the outer shell s2 electron configuration
• First ionisation energy


  The energy needed to remove an electron from each atom in one mole of gaseous atoms
• First ionisation energy


  H(g) → H+(g) + e-
• First and second ionisation energies


  Decrease down the group
• Reactivity of group 2 metals

  Increases down the group
• Second ionisation energy


  The enthalpy change when one mole of gaseous ions with a single positive charge forms one mole of gaseous ions with a double positive charge
• Second ionisation energy


  Ti+(g) → Ti2+(g) + e-
• When the group 2 metals react, they lose their outer shell s2 electrons in redox reactions to form 2+ ions</b>
• Mg will also react slowly with oxygen without a flame
• Mg ribbon will often have a thin layer of magnesium oxide on it formed by reaction with oxygen
• Reaction of Mg with oxygen

  2Mg + O2 → 2MgO
• Mg ribbon needs to be cleaned off by emery paper before doing reactions with Mg ribbon
• If testing for reaction rates with Mg and acid, an un-cleaned Mg ribbon would give a false result because both the Mg and MgO would react but at different rates