periodicity

• periodicity is the recurrence of similar properties of the elements after certain regular intervals when they are arranged in the order of increasing atomic numbers
• s-block is the region of the periodic table where the valence electrons occupy an s-orbital
• p-block is a region of the periodic table where the valence electrons occupy a p-orbital
• d-block is a region of the periodic table where the valence electrons occupy a d-orbital
• 1st ionisation energy is the energy required to remove one mole of the most loosely held electrons from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+
• electronegativity is the ability of an atom to attract a pair of electrons in a covalent bond
• atomic radius is the mean distance from the centre of the nucleus to the boundary of the surrounding shells of electrons
• ionic radius is the radius of an ion in an ionic crystal
• nuclear charge is the combined positive charge of all the protons in the nucleus
• ionic bond is an electrostatic attraction between oppositely charged ions
• covalent bonds are electrostatic attractions between 2 nuclei and the shared pair of electrons between the 2 nuclei
• metallic bonds are electrostatic attractions between delocalised electrons and positive metal ions in a lattice
• trend in atomic size across a period:

  • size of the atom decreases across a period
  • nuclear charge increases
  • shielding stays the same
  • the attraction of the outer electrons to the nucleus increases
  • outer electrons are pulled in closer
• trend in ionisation energy across a period:

  • generally increases across a period
  • decreases between groups 2 and 3 and groups 4 and 5
  • nuclear charge increases, shielding remains the same making the electrons harder to move
• ionisation energy decreases between groups 2 and 3 because in group 3 the outer electron is in a 3p orbital which is at a higher energy level than 3s which is where the outer electron is in group 2. therefore less energy is required to remove the outer electron in group 3
• ionisation energy decreases between group 5 and 6 because in group 6 the outer most electron is paired in the 3p orbital which increases the repulsion in the orbital so the electron is easier to remove
• trends in electronegativity across a period:

  • increases across a period
  • nuclear charge increases, shielding remains the same and atomic radius decreases
  • stronger nuclear attraction on outermost electrons
  • atom will have a larger share of electrons in a covalent bond
• trends in bonding and structure across a period:

  • covalent bonding becomes more common on crossing a period left to right as ionisation energy increases so it becomes more difficult to form metallic structures
  • giant metallic groups 1,2 and 3
  • giant covalent group 4
  • simple molecular groups 5,6 and 7
• trends in melting points across a period(period 3):

  • groups 1,2 and 3 have giant metallic structures and melting point increases with increasing nuclear charge and decreasing atomic radius, melting point increase across the 3
  • group 4 has giant covalent structure, strong covalent bonds increase melting point from group 3 to 4
  • groups 5-7 have simple molecular structures with induced dipole-dipole forces the increased number of electrons the stronger the forces therefore melting point decreases from group 4 to 7 generally
• electrical conductivity increases from sodium to aluminium as the number of delocalised electrons per atom increases. Al3+ has 3 delocalised electrons and Na+ only has 1
• sulfur has the highest melting point in groups 5-7 (simple molecular structure) in period 3 as it exists in S8 molecules which have the highest number of electrons so therefore the strongest induced dipole-dipole forces
• phosphorous exists as P4 molecules, chlorine exists as Cl2 molecules and argon exists as Ar single molecules
• periodicity is the repeating trend of physical and chemical properties across periods of the periodic table
• using periodicity predictions can be made about the likely properties of an element and its compounds
• across each period elements change from metals to non-metals, as you move down the periods this change moves further and further to the right
• the 4s energy level is lower than the 3d energy level, therefore the 4s orbital fills up with electrons before the 3d orbital and the 4s orbital would be emptied of electrons before the 3d orbital during ionisation
• electronic configurations can be shortened by using the electronic configuration of the noble gas that comes before the element on the periodic table and then adding the outer shell portion of the electronic configuration of the element
• first ionisation energy of an element is the energy required to remove one electron from each atom in one mole of the gaseous element to form one mole of gaseous 1+ ions
• the larger the atomic radius, the smaller the nuclear attraction experienced by the outer electrons
• the higher the nuclear charge the large the nuclear attraction on the outer electrons, due to the higher positive charge of the nucleus
• successive ionisation energies are a measure of the amount of energy required to remove each electron in turn
• ionisation energy generally increases across each period
• as a group descends the first ionisation energy decreases as the number of shells increases which increases the atomic radius and the electron shielding increases