Chapter 8

• Group 2 elements


  Also called alkali earth metals, have 2 electrons in their outer shell in the s sub-shell
• Group 2 elements


  • Undergo redox reactions most commonly
  • Oxidise to form 2+ charge with noble gas configuration
• Calcium oxidation


  Ca → Ca2+ + 2e-
• Group 2 elements


  Reducing agents, as they donate 2 electrons to something else
• Down the group


  Reactivity increases
• Reason for increasing reactivity down the group


  • Atomic radius increases, there is more shielding
  • Nuclear attraction decreases, easier to remove outer electron
  • Cations form more easily as atomic radius increases
• Reaction of group 2 metals with oxygen


  2M (s) + O2 (g) → 2MO (s)
• In the reaction with oxygen

  • M = 0 → +2 so the metal has been oxidised (reducing agent)
  • O = 0 → -2 so the oxygen has been reduced (oxidising agent)

  
• Reaction of group 2 metals with water


  M (s) + 2H2O (l) → M(OH)2 (s) + H2 (g)

  
• In the reaction with water

  • M = 0 → +2 so the metal has been oxidised (reducing agent)
  • H = +1 → 0 so the hydrogen has been reduced (oxidising agent)
  • No change in oxidation number for oxygen

  
• Reaction of group 2 metals with acid


  M (s) + 2HCl (aq) → MCl2 (aq) + H2 (g)

  
• In the reaction with acid


  • M = 0 → +2 so the metal has been oxidised (reducing agent)
  • H = +1 → 0 so the hydrogen has been reduced (oxidising agent)
  • No change in oxidation number of chlorine

  
• As you go down group 2, the reactivity increases, the elements become more alkaline and the pH increases when reacted with H2O

  
• Group 2 metals form 2+ ions when oxidised
• The ionisation energy decreases down the group as the attraction between the nucleus and the outer electrons decreases, and the atomic radius increases
• As the total energy input from ionisation energies to form 2+ ions decreases down the group, the elements become more reactive and stronger reducing agents

  
• Group 2 oxides


  React with water, releasing hydroxide ions, to form alkaline solutions

  
• Down the group


  The resulting solutions get more alkaline due to increasing solubility and higher OH- (aq) concentration

  
• Group 2 compounds in agriculture

  • Ca(OH)2 is used to neutralise acidic soils
• Group 2 compounds in medicine


  • Mg(OH)2 is used as a suspension ('milk of magnesia') to neutralise excess acid in the stomach and treat constipation
• Mg(OH)2 is only partially soluble, making the solution slightly alkaline due to low OH- concentration
• Down the group of halogens

  Boiling point increases

  
• Halogens


  • Non-metals and diatomic molecules at room temperature
  • Have a simple molecular structure with weak London forces between molecules
  • The larger the molecule, the stronger the London forces between molecules
• Up the group of halogens


  Reactivity increases
• Colour of halogens


  • Chlorine = very pale green/colourless solution
  • Bromine = yellow/orange solution
  • Iodine = brown solution (sometimes black solid present)

  
• Colour of organic solvent layer with halogens


  • Chlorine = colourless
  • Bromine = yellow
  • Iodine = purple

  
• Colour of halogens alone at room temperature


  • Chlorine = pale green gas
  • Bromine = brown liquid
  • Iodine = grey solid

  
• Reaction of bromide ions and chlorine


  2KBr (aq) + Cl2 (aq) → 2KCl (aq) + Br2 (aq)

  
• In the reaction of bromide ions and chlorine


  • Br = -1 → 0 so the bromine has been oxidised
  • Cl = 0 → -1 so the chlorine has been reduced
  • No change in oxidation number for the potassium

  
• The oxidising power of halogens decreases down the group (they become less reactive)
• Reason for decreasing reactivity down the group


  • Atomic radius increases
  • More electrons in inner shells, more shielding
  • Less nuclear attraction to attract an electron from another species
• Halogens


  Have 7 electrons in their outer shell, 2 in the s sub-shell and 5 in the p sub-shell

  
• Reaction of halogens with metals


  Ca (s) + Cl2 (g) → CaCl2 (s)

  
• Halogens


  Are oxidising agents, gain electrons, causing the metal to oxidise by losing electrons
• Disproportionation reactions are where the same species is both oxidised and reduced
• Reaction of chlorine with water


  Cl2 (g) + H2O (l) → HCl (aq) + HClO (aq)

  
• Reaction of chlorine with cold, dilute aqueous sodium hydroxide

  Cl2 (aq) + 2NaOH (aq) → NaCl (aq) + NaClO (aq) + H2O (l)

  
• In the reaction of chlorine with cold, dilute aqueous sodium hydroxide


  • Chlorine gets oxidised as there is an increase from 0 to +1 in ClO- (aq)
  • Chlorine gets reduced as there is a decrease from 0 to -1 in Cl- (aq)

  
• The mixture of NaCl and NaClO (sodium chlorate(I)) is used as bleach and to disinfect/kill bacteria
• Adding chlorine to water


  Reacts to form the sterilising agents HClO and ClO- (aq) which kill bacteria