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 metals with dilute acid

  • Calcium: very quick
  • Magnesium: quick
  • Zinc: fairly slow
  • Iron: more slow
  • Copper: very slow
• Non-metals hydrogen and carbon are often included in the reactivity series
• Displacement


  A more reactive metal can displace a less reactive metal from a compound
• Gold is very unreactive and is found in the Earth as the metal itself
• Most metals are found as compounds that require chemical reactions to extract the metal
• Reduction


  Involves the loss of oxygen
• Oxidation


  Loss of electrons
• Reduction

  Gain of electrons
• 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
• Determining what has been oxidised and reduced


  1. Look at the changes in the elements

  2. Sodium has lost electrons and been oxidised

  3. Chlorine has not been oxidised or reduced

  4. Hydrogen has gained electrons and been reduced
• Reactions of acids with metals

  1. Acid + metal -> salt + hydrogen

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


  Losing electrons
• Reduction

  Gaining electrons
• Reactions of acids with metals are redox reactions because one substance is oxidised and another is reduced
• Neutralisation of acids and salt production

  1. Acid + alkali -> salt + water

  2. Acid + base -> salt + water

  3. Acid + metal carbonate -> salt + water + carbon dioxide
• Alkali

  Soluble metal hydroxides
• Base

  Insoluble metal hydroxides and metal oxides
• Salt produced from acid + alkali/base


  • Depends on the acid used:
  • Hydrochloric acid (HCl) produces chlorides (XCl)
  • Nitric acid (HNO3) produces nitrates (XNO3)
  • Sulfuric acid (H2SO4) produces sulfates (XSO4)
  • And depends on the positive ion in the base/alkali/carbonate (the metal X)
• The charges on the positive ion from the base/alkali/carbonate and the negative ion from the acid must add up to zero
• Making soluble salts

  Add insoluble solid substance to acid

  2. Keep adding until excess solid sinks to bottom

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

  Measures acidity or alkalinity of a solution

  pH 7 is neutral

  pH < 7 is acidic

  pH > 7 is alkaline
• Acids

  Produce H+ ions in aqueous solutions
• Alkalis

  Produce OH- ions in aqueous solutions
• H+(aq) + OH-(aq) -> H2O(l) is the ionic equation for any neutralisation reaction
• Titration

  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 get more precise results
• Titre


  Volume of acid needed to exactly neutralise the alkali
• 1 dm3 = 1000 cm3
• One mole of a substance in grams is the same as its relative atomic mass in grams
• Titration calculations

  Convert volumes to dm3

  2. Work out moles of NaOH

  3. Use mole ratio from equation to work out moles of HCl

  4. Calculate concentration of HCl
• Strong acid

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


  Partially ionised in aqueous solution (e.g. ethanoic, citric, carbonic acids)
• Strength of acid

  Lower pH (for a given concentration of aqueous solutions)
• As pH decreases by 1 unit, H+ concentration increases by a factor of 10
• Strong/weak and concentrated/dilute are not the same - strong/weak refers to H+ ion concentration, concentrated/dilute refers to amount of substance in a given volume