Periodicity

• s, p, d blocks


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


  • Decreases as you move from left to right across a period
  • Increased number of protons create more positive charge attraction for electrons which are in the same shell
  • Similar shielding
• Periodicity


  A repeating pattern across different periods
• Properties displaying 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
• Atoms of elements in a group


  Have similar outer shell electron configurations, resulting in similar chemical properties
• Elements in periods showing 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 the electrons from the nucleus
  • Shielding of the attraction of the nucleus
• Successive ionisation energies for an element give important information about the electronic structure
• As the number of electrons removed increases

  The ionisation energy increases
• The second ionisation energy of an element is always bigger than the first ionisation energy
• The pattern in the first ionisation energy gives us useful information about electronic structure
• Helium has the largest first ionisation energy
• As you go down a group


  First ionisation energies decrease
• As you go across a period


  First ionisation energies generally increase
• Sodium has a much lower first ionisation energy than Neon
• There is a small drop in first ionisation energy from Mg to Al
• There is a small drop in first ionisation energy from P to S
• 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
• Bonding types


  • Covalent: shared pair of electrons
  • Metallic: electrostatic force of attraction between the metal positive ions and the delocalised electrons
• Structure types


  • Macromolecular: giant molecular structures
  • Giant metallic lattice
• Macromolecular structures


  • Have very high melting points because of strong covalent forces in the giant structure
  • Take a lot of energy to break the many strong covalent bonds
• Giant metallic structures


  • Have high melting and boiling points due to strong electrostatic forces between positive ions and sea of delocalised electrons
  • Malleable as the positive ions in the lattice are all identical, so the planes of ions can slide easily over one another
• Melting and boiling points


  • Na, Mg, Al: Metallic bonding, get stronger the more electrons in outer shell
  • Si: Macromolecular, many strong covalent bonds
  • Cl2, S8, P4: Simple molecular, weak London forces between molecules
  • Ar: Monoatomic, weak London forces
• Similar trend in period 2: Li,Be metallic, B,C macromolecular, N2,O2 molecular, Ne monoatomic
• Historical periodic table


  Ordered according to atomic mass, not proton number
• Mendeleev's periodic table predictions of undiscovered elements turned out to be true
• Modern periodic table


  • Ordered by proton number, not mass number
  • Elements grouped into groups (columns) and periods (rows)
  • Groups relate to number of electrons in outer shell
  • Periods relate to number of electron shells
• Ionisation is endothermic, has a positive value
• Factors affecting ionisation energy


  • Shielding - more electron shells make it easier to remove electrons
  • Atomic size - bigger atoms have weaker attraction between nucleus and outer electrons
  • Nuclear charge - more protons in nucleus increase attraction and require more energy to remove electrons
• Ionisation energy going down a group


  Decreases due to increasing atomic radius and increasing shielding
• Ionisation energy going across a period


  Generally increases due to increasing nuclear charge, with exceptions for aluminium and sulfur
• Aluminium exception


  Outermost electron is in a higher energy sub-shell, slightly further from nucleus and shielded, requiring less energy to remove
• Sulfur exception


  Outermost electrons are paired in the 3p orbital, causing electron repulsion and requiring less energy to remove