Chemrevise

• Atomic radius

  Increases down the Group
• 1st ionisation energy

  • The outermost electrons are held more weakly because they are successively further from the nucleus in additional shells
  • The outer shell electrons become more shielded from the attraction of the nucleus by the repulsive force of inner shell electrons
• Melting points

  • Down the group the melting points decrease
  • The metallic bonding weakens as the atomic size increases
  • The distance between the positive ions and delocalized electrons increases
  • The electrostatic attractive forces between the positive ions and the delocalized electrons weaken
• Group 2 reactions with water


  1. Reactivity of group 2 metals increases down the group

  2. Magnesium reacts in steam to produce magnesium oxide and hydrogen

  3. The other group 2 metals will react with cold water with increasing vigour down the group to form hydroxides

  4. The hydroxides produced make the water alkaline

  5. Fizzing (more vigorous down group)

  6. The metal dissolving (faster down group)

  7. The solution heating up (more down group)

  8. With calcium a white precipitate appearing (less precipitate forms down group)
• Reactions of group 2 metals with oxygen

  1. The group 2 metals will burn in oxygen

  2. Magnesium burns with a bright white flame

  3. Magnesium will also react with warm water, giving a different magnesium hydroxide product
• Magnesium oxide


  A white solid with a high melting point due to its ionic bonding
• Reactions of group 2 metals with chlorine

  The group 2 metals will react with chlorine
• • The reactivity increases down the group as the atomic radii increase there is more shielding
  • The nuclear attraction decreases and it is easier to remove (outer) electrons and so cations form more easily
• Reactions of the oxides of group 2 elements with water


  1. The ionic oxides are basic as the oxide ions accept protons to become hydroxide ions in this reaction (acting as a bronsted lowry base)

  2. MgO (s) + H2O (l) Mg(OH)2 (s)

  3. pH 9

  4. Mg(OH)2 is only slightly soluble in water so fewer free OH- ions are produced and so lower pH
• Reactions of the oxides of group 2 elements with acids

  1. MgO (s) + 2 HCl (aq) MgCl2 (aq) + H2O (l)

  2. 2HNO3 (aq) + Mg(OH)2 (aq) Mg(NO3)2 (aq)+ 2H2O (l)
• Reactions of the hydroxides of group 2 elements with acids

  2HCl (aq) + Mg(OH)2 (aq) MgCl2 (aq)+ 2H2O (l)
• Solubility of group 2 sulfates

  • Group II sulfates become less soluble down the group
  • BaSO4 is the least soluble
  • If Barium metal is reacted with sufuric acid it will only react slowly as the insoluble barium sulfate produced will cover the surface of the metal and act as a barrier to further attack
• Solubility of group 2 hydroxides

  • Group II hydroxides become more soluble down the group
  • All Group II hydroxides when not soluble appear as white precipitates
  • Calcium hydroxide is reasonably soluble in water
  • Barium hydroxide would easily dissolve in water
  • Magnesium hydroxide is classed as insoluble in water
• Thermal decomposition of group 2 carbonates

  1. The ease of thermal decomposition decreases down the group

  2. CaCO3(s) CaO(s) + CO2(g)

  3. MgCO3(s) MgO(s) + CO2(g)

  4. Group 2 carbonates become more thermally stable going down the group

  5. As the cations get bigger they have less of a polarising effect and distort the carbonate ion less

  6. The C-O bond is weakened less so it less easily breaks down
• Group 1 carbonates do not decompose with the exception of lithium
• Thermal decomposition of group 2 nitrates (V)


  1. Group 2 nitrates decompose on heating to produce group 2 oxides, oxygen and nitrogen dioxide gas

  2. The ease of thermal decomposition decreases down the group

  3. The explanation for change in thermal stability is the same as for carbonates
• Group 1 nitrates, with the exception of lithium nitrate, do not decompose in the same way as group 2 nitrates
• Flame tests

  • Lithium: Scarlet red
  • Sodium: Yellow
  • Potassium: lilac
  • Rubidium: red
  • Caesium: blue
  • Magnesium: no flame colour (energy emitted of a wavelength outside visible spectrum)
  • Calcium: brick red
  • Strontium: red
  • Barium: apple green
• The oxidation reactions of halide ions by halogens

  1. Chlorine will displace both bromide and iodide ions; bromine will displace iodide ions

  2. Cl2(aq) + 2Br – (aq) 2Cl – (aq) + Br2(aq)

  3. Cl2(aq) + 2I – (aq) 2Cl – (aq) + I2(aq)

  4. Br2(aq) + 2I – (aq) 2Br – (aq) + I2(aq)
• Observations of halogen reactions

  • Chlorine = very pale green solution (often colourless)
  • Bromine = yellow solution
  • Iodine = brown solution (sometimes black solid present)
• Observations if an organic solvent is added

  1. Chlorine = colourless

  2. Bromine = yellow

  3. Iodine = purple
• Halogens


  • Fluorine (F2): very pale yellow gas, highly reactive
  • Chlorine (Cl2): greenish, reactive gas, poisonous in high concentrations
  • Bromine (Br2): red liquid, gives off dense brown/orange poisonous fumes
  • Iodine (I2): shiny grey solid sublimes to purple gas
• Disproportionation reactions of chlorine

  1. Cl2(g) + H2O(l) HClO(aq) + HCl (aq)

  2. Chlorine is both simultaneously reducing and oxidising changing its oxidation number from 0 to -1 and 0 to +1

  3. If some universal indicator is added to the solution it will first turn red due to the acidity of both reaction products

  4. It will then turn colourless as the HClO bleaches the colour
• Chlorine is used in water treatment to kill bacteria
• Covalent bond


  Bond formed by the sharing of a pair of electrons between two atoms
• As one goes down the group

  The electronegativity of the elements decreases
• As one goes down the group

  The atomic radii increases due to the increasing number of shells
• The reactivity of the halogens decreases down the group


  As the atoms get bigger with more shielding, they less easily attract and accept electrons
• Disproportionation


  A reaction where an element simultaneously oxidises and reduces
• Chlorine with water


  Cl2(g) + H2O(l) → HClO(aq) + HCl(aq)
• Chlorine is both simultaneously reducing and oxidising changing its oxidation number from 0 to -1 and 0 to +1
• If some universal indicator is added to the solution it will first turn red due to the acidity of both reaction products. It will then turn colourless as the HClO bleaches the colour.
• Chlorine


  Used in water treatment to kill bacteria
• The benefits to health of water treatment by chlorine outweigh its toxic effects
• Reaction of halogens with cold dilute NaOH solution


  Cl2(aq) + 2NaOH(aq) → NaCl(aq) + NaClO(aq) + H2O(l)
• The mixture of NaCl and NaClO is used as Bleach and to disinfect/ kill bacteria
• Reaction of 3I2(aq) with 6OH-(aq)

  5I-(aq) + IO3-(aq) + 3H2O(l)
• IUPAC convention


  The various forms of sulfur and chlorine compounds where oxygen is combined are all called sulfates and chlorates with relevant oxidation number given in roman numerals
• Reaction of halogens with hot dilute NaOH solution

  3Cl2(aq) + 6NaOH(aq) → 5NaCl(aq) + NaClO3(aq) + 3H2O(l)

  3I2(aq) + 6NaOH(aq) → 5NaI(aq) + NaIO3(aq) + 3H2O(l)
• Oxidation reactions of metals and metal ions by halogens


  3Cl2(g) + 2Fe(s) → 2FeCl3(s)

  Br2(l) + 2Na(s) → 2NaBr(s)

  Cl2(g) + 2Fe2+(aq) → 2Cl-(aq) + 2Fe3+(aq)

  2I-(aq) + 2Fe3+(aq) → I2(aq) + 2Fe2+(aq)

  Br2(l) + Mg(s) → MgBr2(s)