Chemical changes

• Metal oxides


  Metals + oxygen -> metal oxides
• Oxidation


  Gain of oxygen
• Reduction


  Loss of oxygen
• Reactivity series


  • When metals react with other substances, metal atoms form positive ions
  • Reactivity of a metal is related to its tendency to form positive ions
  • Metals can be arranged in order of their reactivity in a reactivity series
• Metals in order of reactivity

  • Potassium
  • Sodium
  • Lithium
  • Calcium
  • Magnesium
  • Zinc
  • Iron
  • Copper
• Reactions of metals with water

  • Potassium: violent
  • Sodium: very quick
  • Lithium: quick
  • Calcium: more slow
• Reactions of acids with metals

  1. Acid + metal -> salt + hydrogen

  2. These are redox reactions - one substance is reduced and another is oxidised

  3. Identify which substances are oxidised and reduced by looking at electrons gained and lost (following OIL RIG)
• Reactions of metals with dilute acid

  • Calcium: very quick
  • Magnesium: quick
  • Zinc: fairly slow
  • Iron: more slow
  • Copper: very slow
• Redox reaction


  A reaction where one substance is reduced and another is oxidised
• Electrolysis


  The process of breaking down an ionic substance into its elements by passing an electric current through it
• Non-metals hydrogen and carbon are often included in the reactivity series
• Neutralisation of acids and salt production

  1. Acid + alkali -> salt + water

  2. Acid + base -> salt + water

  3. Acid + metal carbonate -> salt + water + carbon dioxide
• The process of electrolysis

  1. Ionic substance is melted or dissolved

  2. Ions are free to move about

  3. Current is passed through the molten or dissolved substance

  4. Substance is broken down into its elements
• Displacement


  A more reactive metal can displace a less reactive metal from a compound
• Cathode


  The negative electrode where positively charged ions move to
• Gold is very unreactive and is found in the Earth as the metal itself
• Salt produced in alkali and base reactions

  • Depends on the acid used:
  • Hydrochloric acid (HCl) produces chlorides (XCl)
  • Nitric acid (HNO3) produces nitrates (XNO3)
  • Sulfuric acid (H2SO4) produces sulfates (XSO4)
  • Also depends on the positive ion in the base, alkali or carbonate (the metal X)
• Anode


  The positive electrode where negatively charged ions move to
• Most metals are found as compounds that require chemical reactions to extract the metal
• The charges on the positive ion from the base/alkali/carbonate and the negative ion from the acid must add up to zero
• Electrolysis of molten ionic compounds

  1. Metal is produced at the cathode

  2. Non-metal is produced at the anode
• Reduction


  Involves the loss of oxygen
• Making soluble salts


  1. Add the chosen solid insoluble substance to the acid, the solid will dissolve

  2. Keep adding until excess solid sinks to the bottom, indicating the acid has been neutralised

  3. Filter out excess solid, evaporate some water, then leave the rest to evaporate slowly (crystallisation)
• pH scale


  • Measures the acidity or alkalinity of a solution
  • pH 7 is neutral
  • pH < 7 is acidic
  • pH > 7 is alkaline
• Electrolysis of molten ionic compounds

  • Metal is the positive ions
  • Non-metal is the negative ions
• Neutralisation reaction


  H+(aq) + OH-(aq) -> H2O(l)
• Using electrolysis to extract metals

  1. Metals more reactive than carbon are extracted by electrolysis of molten compounds

  2. Large amounts of energy are used to melt the compounds and produce the electrical current

  3. Aluminium is manufactured by electrolysis of a molten mixture of aluminium oxide and cryolite using carbon as the positive electrode
• Oxidation


  Loss of electrons
• Electrolysis of aluminium


  • Aluminium oxide has a very high melting point, so it is mixed with cryolite to lower the melting point
  • The positive electrodes need to be continually replaced because oxygen is formed, which reacts with the carbon forming carbon dioxide
• Reduction


  Gain of electrons
• Titration


  1. Wash burette with dilute HCl and water

  2. Fill burette to 100cm3 with acid

  3. Use 25cm3 pipette to add 25cm3 of alkali to conical flask

  4. Add indicator (e.g. phenolphthalein)

  5. Add acid from burette to alkali until end-point (as shown by indicator)

  6. Titre is difference between first and second burette readings

  7. Repeat to gain more precise results
• Writing ionic equations


  1. If sodium is oxidised, it has lost an electron, leaving it with a +1 charge, so the ionic equation is: Na -> Na+ + e-

  2. If sodium +1 ion is reduced, it has gained an electron, leaving it with a charge of zero, so the ionic equation is: Na+ + e- -> Na

  3. The charges on each side of the equation should add up to the same number
• Metals that react with carbon can also be extracted by electrolysis
• Electrolysis of aqueous solutions

  1. Ions discharged depend on the relative reactivity of the elements

  2. At the cathode, hydrogen is produced unless the metal is less reactive than hydrogen

  3. At the anode, if OH- and halide ions are present, one of the halide ions will be produced, otherwise oxygen is formed
• Titration calculations


  • 1dm3 = 1000cm3
  • One mole of a substance in grams is the same as its relative atomic mass in grams
• Determining what has been oxidised and reduced


  1. Look at the changes in the elements involved

  2. e.g. 2Na + 2HCl -> 2NaCl + H2

  Sodium has lost electrons and been oxidised (2Na -> 2Na+ + 2e-)

  Hydrogen has gained electrons and been reduced (2H+ + 2e- -> H2)
• In aqueous solutions, water molecules break down producing H+ and OH- ions that are discharged
• Example titration calculation: 25cm3 of dilute HCl is neutralised by 20cm3 of 0.5 mol/dm3 NaOH. What is the concentration of the HCl?
• Half equation


  Representation of the reaction at an electrode, where the small number is the same as the 2 larger numbers, and electrons are represented by 'e-'
• Strong acid


  Completely ionised in aqueous solution (e.g. HCl, HNO3, H2SO4)