Reactivity trends

• Group 2 elements:

  • Lose 2 electrons easily makes it a good reducing agent
  • Ca —> Ca2+ + 2e-
  • redox reactions with oxygen + water
  • 2Ca + O2 —> 2CaO
  • Ca + 2H2O —> Ca(OH)2 + H2
• Ionisation energy & reactivity:

  • More reactive down a group (size increasing, shielding increases)
  • Ionisation energy decreases down group (increasing size and shielding)
• Group 2 compounds:

  • Oxides CaO + H2O —> Ca(OH)2 —> Hydroxides not very soluble
  • Solubility increases down the group
  • Ca(OH)2 used to control pH of soil
  • Mg(OH)2 used as an antacid (indigestion tablet)
  • Mg(OH)2 + 2HCl —> MgCl2 + 2H2O
• • Halogens are all diatomic non-polar molecules, consisting of a single covalent bond
  • All react to form -1 anions, as they all have 7 outer shell electrons, so are powerful oxidising agents
  • Their oxidation states varies in compounds with oxygen present, with the exception o fluorine which is always -1
  • X2 + 2e- —> 2X-
• Halogen physical properties:

  • The boiling points of the halogens increase down the group
  • Flourine & Chlorine are gases at RTP
  • Bromine is a liquid at RTP
  • Iodine is solid at RTP
  • The increase in atomic radius due to more electron shells, means there are more electrons so Greater London forces
• Halogen reactivity:

  • Halogens have to gain electrons so the reactivity decreases down the group
  • The increase in nuclear charge down the group would increase the ability to attract an electron to fill the outer shell
  • However the increased radius & thus increased electron shielding makes it harder to attract an electron, so the reactivity decreases
• Displacement reactions:

  • A2 + 2KB —> 2KA+ B2
  • Place a potassium halide to a height of 2cm into a test tube
  • Add an equal volume of a halogen solution
  • Stopper the test tube & shake carefully
  • Add an equal volume of cyclohexane
  • Stopper the test tube & shake carefully
  • Observe the colour of the organic layer at the top
• Anion Tests:

  • Carbonates
  • Sulfate
  • Halide
• • Carbonate ions (CO3 2-) are identified by adding hydrochloric acid, which causes them to effervesce
  • CO3 2- + 2H+ (aq ) —> H2O (l) + CO2 (g)
  • The gas can be passed through limewater as a confirmatory test, as the CO2 formed will form a precipitate of CaCO3.
• • Sulfate ions (SO4 2-) are identified by adding hydrochloric acid then barium chloride or nitrate solutions to form a precipitate of barium sulfate.
  • Ba 2+ (aq) + SO4 2- (aq) —> BaSO4 (s)
• Halide ions (Cl-, Br-, & I- ) can be identified from positive tests with acidified silver nitrate:

  • Chlorine —> white precipitate that dissolves in NH3
  • Bromine —> cream precipitate that dissolves in concentration NH3
  • Iodine —> yellow precipitate that is insoluble in NH3
• Cation test:

  • Ammonium ions can be identified by warming them with NaOH to generate ammonia gas
  • NH4+ (aq) + OH- (aq) —> NH3(g) + H2O(l)
  • The gas can then be identified by putting damp red litmus paper into the vapour, which will turn blue with NH3.
• Uses of chlorine:

  • Often used in water treatment as it reacts with water to form solutions that kill bacteria
  • But there are concerns about chlorinated hydrocarbons forming
  • Cl2 (aq) + H2O (l) —> HCl (aq) + HClO (aq)
  • Chlorine also reacts with sodium hydroxide to form NaCl, water and a variety of chlorides (NaClOx) .i.e. bleach
  • Cl2 (g) + 2NaOH (aq) —-> NaCl (aq) + NaClO (aq) + H2O (l)
  • Both of these are examples of disproportionation reactions as the chlorine is both oxidised and reduced in the reaction
• First ionisation energy decreases between group 2 and 3:

  • Decrease between 2 to 3 because in group 3 the outermost electrons are in p orbitals whereas in group 2 they are in s orbital, so the electrons are easier to be removed
• First ionisation decreases between 5 to 6:

  • decrease between 5 to 6 is due to the group 5 electrons in p-orbital are single electrons and in group 6 the outermost electrons are spin paired, with some repulsion. Therefore the electrons are slightly easier to remove
• There is a decrease in ionisation energy between the end of one period and the start of next as there is an increase in atomic radius and electron shielding
• First ionisation energy decreases down a group:

  • Shielding increases --> weaker attraction
  • Atomic radius increases --> distance between the outer electrons and nucleus increases --> weaker attraction
  • Increase in number of protons is outweighed by increase in distance and shielding
• Group 2 metal properties:

  • High melting and boiling points
  • Low density metals
  • Form colourless white compounds
• First ionisation energy decreases down group 2:

  • Number of filled electron shells increases down the group --> increasing shielding
  • Increased atomic radius --> weaker force between outer electron and nucleus --> less energy needed to remove electron
• Beryllium doesn't react with water
• Magnesium reacts very slowly with water
• Beryllium oxide is insoluble in water
• Hydroxide solubility:

  • increases down group 2
  • magnesium hydroxide is slightly soluble
  • barium hydroxide creates a strong alkaline solution
• Calcium hydroxide is used to neutralise soil
• Magnesium hydroxide ( milk of magnesia) --> antacid to treat indigestion, heartburn
• Calcium carbonate is present in limestone and marble and used in building construction
• Chlorine is pale green in water
• Bromine is orange in water
• Iodine is brown in water
• Chlorine is pale green in cyclohexane
• Bromine is orange in cyclohexane
• Iodine is violet in cyclohexane