Periodicity notes

• Ionisation energy


  The energy required to remove one electron from each atom in a mole of gaseous atoms to form one mole of gaseous 1+ ions
• Ionisation energy


  • Always measured in the gaseous state
  • 1 electron is removed at a time
  • 1 electron is removed from each atom in a mole of atoms/1 electron from each ion in a mole of ions
• Ionisation requires energy
• Subsequent ionisations


  Take more energy as there are now less electrons in ratio to the number of protons, so there is less repulsion between the electrons, and a greater attractive force per electron to the nucleus
• Subsequent ionisations


  Take more energy as the less electrons there are, the less shells there are, and so the less shielding there is
• 1st electron removed


  Called first ionisation energy
• 1st ionisation energy


  X(g) -> X+(g) + e-
• 2nd ionisation energy


  X+(g) -> X2+(g) + e-
• For each successive ionisation energy, the proton:electron ratio increases, therefore there is more positive charge attracting the remaining electrons each time
• This causes the outermost electron (the one being removed) to be attracted most strongly
• Electrons in the inner shells experience much greater attraction from the protons in the nucleus due to shorter distance between electrons and the nucleus, and less shielding from inner shells
• There are big jumps in ionisation energy when the first electron is taken from a new shell as that shell (and all electrons in it) have a much stronger attraction to the nucleus, so require more energy to remove
• The number of electrons before the first big jump is the number of electrons in the outermost shell of an element, and therefore which group it is in
• Periodicity


  The regular repeating patterns in the properties (chemical- how they react and physical- boiling points, metal or non-metal) of elements
• The groups show elements with similar properties, the periods show a trend (gradual change) in properties
• Factors affecting first ionisation energy

  • Nuclear charge (how many protons are in the nucleus)
  • Atomic radius (distance between the nucleus and outermost electron)
  • Shielding (the more shells there are between the outermost electron and the nucleus, the weaker the attraction and the less energy needed for ionisation)
• First ionisation energy decreases down the group


  As increased nuclear charge is outweighed by increased atomic radius and shielding
• First ionisation energy increases left to right across a period


  Because the nuclear charge is increasing, shielding is the same as all in same period, so greater pull from the protons in the nucleus on the outermost electrons, decreasing the atomic radius (electrons pulled closer), so more energy will be required to remove the outermost electron as it is more strongly held
• Boron dips in first ionisation energy because we are removing an electron from a higher sub level, as the 2p sublevel is higher in energy than the 2s, making the electron easier to remove
• Oxygen dips in first ionisation energy because it is the first in the period to have a paired electron in its 2p shell, and this paired electron will experience more repulsion than if it was unpaired, making it slightly easier to remove
• Aluminium dips in first ionisation energy because we are removing the first electron for a higher subshell (3p rather than 3s), making the electron easier to remove
• Sulphur dips in first ionisation energy because it is the first in the period to have a paired electron in the 3p shell, so this electron experiences slightly more repulsion than if it was unpaired, causing it to be easier to remove
• Melting points of period 3 elements

  • The first three elements are metals with high melting and boiling points due to strong forces of attraction between positive metal ions and delocalised electrons
  • Melting points increase across the period for these three as the ionic charge is greater, and the ionic radius is therefore smaller because the electrons are more strongly attracted to the nucleus
  • There is a rapid increase in melting point at silicon as it is a non-metal with a giant covalent structure, so with strong covalent bonds
  • The remaining four elements have simple molecular structure, and melting point at this point gets lower again as the weak intermolecular forces between molecules (van der waals, hydrogen bonding..etc) require very little energy to break
• Period 2 shows an almost identical trend in melting points of the elements as period 3