Elements of Group 1 and 2

• down group 2, atomic radius increases. This is because:

  • nuclear charge increases.
  • this increase is cancelled out by the increase in shielding.
  • atoms have more shells of electrons.
• down group 2, melting point decreases. This is because:

  • metallic bonding weakens as the ionic radius of the positive metal ion increases.
  • the charge density of the metal ion thus decreases.
  • thus there's a weaker electrostatic attraction between ions and delocalised electrons, which weakens metallic bonds.

  the exception is from magnesium to calcium, as there is an increase in melting point due to magnesium having a different structure.
• down group 2, the first ionisation energy of elements decreases. This is because:

  • the atomic radius increases.
  • thus outermost electrons are held successively further from the nucleus.
  • thus outer electrons become more shielded from attraction of the nucleus by repulsion of inner electrons.
  • thus its easier for outer electrons to be lost.
• down group 2, reactivity increases. This is because:

  • these metals react by losing their two outer most electrons.
  • as atomic radius increases, outermost electrons are held successively further from the nucleus.
  • outer electrons become more shielded from attraction of the nucleus by core electrons.
  • thus 2 outermost electrons are more readily lost.
• group 2 metals react with oxygen to form metal oxides

  e.g. 2M (s) + O2 (g) = 2MO (s)

  
  

  In addition:

  • strontium and barium can also make metal peroxides (MO2 (s))
  • magnesium reacts with oxygen in a bright white flame
• group 2 metals react with chlorine to form metal chlorides

  e.g. M (s) + Cl2 (g) = MCl2
• magnesium reacts with steam to produce magnesium oxide and hydrogen gas in a bright white flame.

  e.g. Mg (s) + H2O (g) = MgO (s) + H2 (g)
• magnesium reacts with warm water to produce magnesium hydroxide:

  e.g. Mg (s) + 2 H2O (l) = Mg(OH)2 (s) + H2 (g)
• group 2 metals react with cold water with increasing vigour down the group, making M(OH)2 (aq) and H2 (g)

  
  

  they react with increasing vigour because their first and second ionisation energies decrease down group 2. Hence forming the M2+ ion requires less energy.
• group 2 metal oxides react with water to make aqueous metal hydroxides only
• group 2 metal oxides react with dilute acids to form aqueous salts and water (they're basic)
• group 2 metal hydroxides react with dilute acids to make aqueous salts and water.
• group 2 hydroxides become more soluble down the group.

  e.g. Mg(OH)2 is the only one thats completely insoluble (will form when Mg2+ ions are reacted with sodium hydroxide solution)
• group 2 sulphates become less soluble down the group.

  e.g. BaSO4 is the only one that completely insoluble (will be formed when Ba2+ ions are reacted with sulfuric acid)
• flame test - metal ions producing characteristic flame colours when heated.
• flame test procedure:

  1. clean a nichrome wire by dipping it in HCL and heating it in a roaring bunsen flame until no colour is produced by the wire.
  2. moisten the wire with HCL, and dip it into the sample (dissolved in HCL).
  3. heat the sample in a roaring bunsen flame

  nichrome wires are used because they have a high mp, they're unreactive, and they dont produce a characteristic flame colour themselves.
• Flame Colours:

  Li+ = red

  Na+ = yellow

  K+ = lilac

  Rb+ = red

  Cs+ = blue-violet

  Be2+ = no colour

  Mg2+ = no colour

  Ca2+ = brick red

  Sr2+ = red

  Ba2+ = apple-green
• How flame colour arises:

  • thermal energy excites an electron to a higher energy orbital.
  • the electron drops from the excited state back to the ground state.
  • the drop of an electron releases energy corresponding to a point on the visible light spectrum.
  • if no colour is observed, the energy is released outside of the visible region of the electromagnetic spectrum.
• group 1 nitrates (except lithium) decompose to make metal nitrites (s) (MNO2) and oxygen.
• lithium nitrate + group 2 nitrates decompose to form a metal oxide (s), nitrogen dioxide, and oxygen.

  
  

  * G1 and G2 nitrates decompose to different products because of the difference in charge density, and hence polarisation.
• group 1 carbonates decompose to make metal oxides (s) and carbon dioxide.

  
  

  (lithium carbonate is the only carbonate that can decompose in a bunsen flame — behaves like a G2 carbonate. This is because the low charge of the G1 cation hardly polarises the carbonate anion)
• group 2 carbonates decompose to make metal oxides (s) and carbon dioxide.

  
  

  * G1 and G2 carbonates decompose to different products because of the difference in charge density, and hence polarisation.
• for group 1 and group 2 nitrates and carbonates, thermal stability increases down the group.
• thermal stability increasing down the group explanation:

  • down the group, the ionic radius of the metal ion increases.
  • the metal ion thus becomes less charge dense.
  • this polarises the anion less, thus the bonds (C-O or N-O) in the anion are weakened less.
  • thus its harder for the anion to decompose down the group.
• small highly charged ions (more charge dense) polarise (pull electron density towards themselves) anions more.
• metal + acid = salt + hydrogen