Periodicity

• Ionisation energy - the minimum amount of energy required to remove 1 mole of electrons from 1 mole of gaseous atoms.
• Ionisation is always an endothermic process as it requires energy. Value is always positive.
• Shielding - the energy required depends on how many electron shells there are in between the positive nucleus and the negative electron that's being removed. If there's more electron shells, less energy is required as there's a weaker attraction.
• An increased atomic size means that the negative electron is going to be further away from the positive electron, so the attractive forces reduce, making it easier to remove electrons.
• Nuclear charge - the amount of protons in an atom's nucleus.
• The more protons there are in the nucleus, the larger the nuclear charge -> the bigger the attraction between the nucleus and the outer electrons -> the more energy required to remove the electrons.
• As you go down a group...

  • ionisation energy decreases
  • atomic radius increases - outer electrons further away from the nucleus so weaker attraction so energy required to remove electrons decrease
  • shielding increases - more shells in between electron and nucleus so weaker attraction & less energy required to remove electrons.
• As you go across the period...

  • ionisation energy increases - more energy is required to remove the electron
  • nuclear charge increases - higher nuclear attraction
  • similar shielding - distance between electrons and nucleus decreases slightly.
• Successive ionisation - the removal of more than 1 electron from the same atom.
• Graphite:

  • each carbon bonded 3 times; 4th is delocalised
  • very high melting point due to lots of strong covalent bonds
  • layers slide easily due to the weak forces in between
  • conduct electricity due to presence of delocalised electrons
  • layers are far apart compared to usual covalent bond length -> low density/light (ideal for pencils)
  • insoluble as covalent bonds are hard to break.
• Diamond:

  • each carbon is bonded 4 times; giving the tetrahedral shape
  • tightly packed -> rigid arrangement so it's a good heat conductor
  • can be cut into gemstones (rings, necklaces, etc.)
  • very high melting point due to lots of strong covalent bonds
  • doesn't conduct electricity due to no delocalised electrons
  • insoluble due to strong covalent bonds that are hard to break.
• Graphene:

  • lightweight & transparent
  • excellent electricity conductor (delocalised electrons)
  • really strong (delocalised electrons strengthen the bond)
• Metallic bonding - The electrostatic attraction between the positive metal ions and the delocalised electrons.
• The more electrons an atom can donate to the delocalised system, the higher the melting point.

  e.g. Mg can donate 2 electrons whereas Na can only donate 1, so Mg has a higher melting point than Na.
• Metals are good thermal conductors as the delocalised electrons can transfer kinetic energy.
• Metals are good electrical conductors as they have mobile delocalised electrons that can carry a current.
• Metals have high melting points as they require a lot of energy to break the bonds.
• Metals are insoluble as metallic bonds are too strong to break.
• The melting point of period 3 elements increases since metallic bonds get stronger.
• Silicon has a giant covalent structure (same as diamond), so it has the highest melting point in period 3 as it requires a lot of energy to break the strong bonds.
• Across period 3, the melting point of the metals increases as...

  • there's an increased positive charge (the ions; e.g. Na = +1, Mg = +2, Al = +3)
  • increased number of delocalised electrons
  • smaller ionic radius
• Phosphorus (P4) has a simpler molecular structure than silicon, so it has a lower melting point. In this case, the melting point is determined by the strength of the London forces, not the strength of the covalent bonds.
• Sulfur (S8) has a higher melting point than phosphorus due to its larger simpler molecular structure. Since there are 8 sulfur atoms (phosphorus only has 4 of its atoms), it has stronger London forces, therefore requires more energy to break the forces.
• Chlorine is a diatomic molecule that has a smaller simple molecular structure than sulfur. It has weaker London forces than sulfur, so lower melting point.
• Argon is a monoatomic molecule so it has weak London forces, therefore a lower melting point compared to other elements in the period.
• The different types of bonding include:

  • giant covalent bonding
  • simple covalent bonding
  • giant ionic bonding
  • metallic bonding.
• In simple covalent molecules, some molecules will dissolve in water, some won't. It depends on the polarity of the molecule. Polar = will dissolve. Non-polar = won't dissolve.
• Polar molecules dissolve well in polar solvents like water. Non-polar molecules don't.