Inorganic Chemistry

• Alkali Metals


  • Have one electron in their outer shell
  • Have low density
  • Are stored under oil (to prevent reactions with oxygen or water)
  • Are soft (can be cut with knife)
• How Group 1 elements react with non-metals

  Form ionic compounds which are soluble white solids which form colourless solutions
• Group 1 elements all have one electron in their outer shell
• Reactivity of Group 1 elements

  Increases as the atoms get larger and the distance between the nucleus and the outer electrons increases, allowing them to more easily lose electrons
• Group 7


  • Have 7 electrons in outer shell
  • Form coloured vapours
  • Form diatomic molecules
  • Form ionic salts with metals
  • Form molecular compounds with non-metals
• Group 7 elements and their states of matter

  • Fluorine, F. F2 is a pale yellow gas.
  • Chlorine, Cl. Cl2 is a pale green gas.
  • Bromine, Br. Br2 is dark brown liquid (gives off orange vapour at room temperature).
  • Iodine, I. I2 is a grey solid (gives off purple vapour when heated).
  • Astatine, At. At2 solid at room temperature.
• Changes in Group 7 moving down the group

  • Higher relative molecular mass
  • Higher melting and boiling point
  • Less reactive – electrons less easily gained
• A more reactive halogen displacing a less reactive one from an aqueous solution of its salt


  1. Cl2 + 2 NaBr → Br2 + 2 NaCl, or Cl2 + 2Br– → Br2 + 2 Cl–; orange colour of Br2 appears

  2. Cl2 + 2 NaI → I2 + 2 NaCl, or Cl2 + 2I– → I2 + 2 Cl–

  Br2 + 2 NaI → I2 + 2 NaBr, or Br2 + 2I– → I2+ 2 Br–; brown colour of I2 appears
• Reactivity decreases down Group 7 as the atoms get larger and an incoming electron will be less tightly held by the attractive forces from the nucleus
• Transition metals


  • Form ions with different charges
  • Form coloured compounds
  • Have catalytic properties
• Solubility rules for salts

  • K+, Na+,NH4+ - all salts soluble

  NO3- - all salts soluble

  SO42− - all soluble, except Pb2+, Ba2+, Ca2+

  Cl− - all soluble, except Pb2+, Ag+

  CO32− - all INsoluble, except K+, Na+,NH4+
• Rusting


  Process of forming hydrated iron (III) oxide when iron is exposed to water and air, leaving a brown deposit
• Preparing a sample of PbSO4


  Mix a solution of a soluble Pb (II) salt, e.g. nitrate, with a source of sulfates, e.g. Na2SO4

  PbSO4 is insoluble in water - precipitate forms. Filter to collect the precipitate. Wash with H2O. Leave to dry.
• Preparing a sample of NaCl from NaOH and HCl

  Add one reagent by burette until indicator shows neutralisation, then boil to remove solvent and leave to crystallise. Filter and dry the product.
• Checking identity of a pure liquid

  Measure the boiling point
• Showing a liquid contains pure water

  Add the liquid to anhydrous CuSO4 crystals - formation of blue, hydrated copper (II) sulfate indicates presence of water
• Oxidising agent


  Species that gets reduced in a redox reaction, causing oxidation of another substance
• Reducing agent


  Species that gets oxidised in a redox reaction, causing reduction of another substance
• Testing for presence of NH4+ ions

  Add NaOH, shake gently, put damp litmus paper near outlet - blue litmus indicates NH3 formation from NH4+
• Oxidising agent


  A species that gets reduced in a redox reaction (gains electrons; causes the oxidation of another substance)
• Reducing agent


  A species that gets oxidised in a redox reaction (loses electrons; causes the reduction of another substance)
• How to test for the presence of NH4+ ions

  1. Add some NaOH to the aqueous solution of the tested salt

  2. Shake gently

  3. Put a damp litmus paper near the outlet of the test tube

  4. The damp litmus paper will turn blue if NH4+ were present in your solution (NH3 formation)
• Colours of flames observed when lithium, sodium, and potassium burn in oxygen

  • Crimson-red, Li
  • Yellow-orange, Na
  • Lilac, K
• How to conduct a titration

  1. Rinse the pipette with a solution of unknown concentration, use the pipette to measure out the known volume of this solution

  2. Add an indicator (a substance that changes colour at the end of titration)

  3. Rinse the burette with a solution of known concentration, discard the liquid, use a burette to gradually add the solution of a known concentration

  4. When indicator changes colour (at the end point), the volume added is recorded

  5. It is important to get concordant volume results - they have to lie close to each other

  6. Suitable calculations are performed to find the concentration
• Oxidation


  The addition of oxygen to a substance
• Reduction


  The loss of oxygen from a substance
• Reactivity series of metals

  • Shows the metals in order of their reactivity
  • Metals above H2 in reactivity series react with acid to produce H2, the more reactive the metal the quicker and more violent reaction with acid
  • Metals below H2 don't react with acids
  • Not all metals above H2 react with water - mostly Group I and II metals, Aluminium is the borderline case
• Displacement reaction


  A reaction where a more reactive metal displaces a less reactive metal in a compound
• How unreactive metals are found in Earth

  In their natural state, that being the metal alone rather than in a compound
• How metals less reactive than carbon are extracted

  Reduction with carbon, carbon displaces the metal in a metal oxide - gets oxidised to carbon oxides, metal from the metal oxide gets reduced to the pure metal
• How metals more reactive than carbon are extracted

  By electrolysis
• Oxidation


  Loss of electrons
• Reduction


  Gain of electrons
• Reaction between metals and acids

  Metal + acid → salt + hydrogen, Redox reaction, also a displacement reaction
• Metals in the reactivity series that will react with acid

  • Those above hydrogen
• Neutralisation reaction


  Acid + Base → Salt + Water
• Reaction between metal carbonate and acid

  Metal Carbonate + Acid → Salt + Water + Carbon Dioxide
• Reaction between metal oxide and acid

  Metal Oxide + Acid → Salt + Water
• Redox reaction


  A reaction where both oxidation and reduction occurs
• When magnesium reacts with hydrochloric acid

  Magnesium has lost electrons and thus has been oxidised (Mg to Mg2+), the hydrogen in HCl has gained electrons and thus has been reduced (H+ to H2)