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
• Extraction of metals


  • Gold is very unreactive and found in the Earth as the metal itself
  • Most metals are found as compounds that require chemical reactions to extract the metal
  • Metals less reactive than carbon can be extracted from their oxides by reduction with carbon
• 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
• Identifying oxidation and reduction in an equation

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

  2. HCl is made up of H+ and Cl- ions & NaCl is made up of Na+ and Cl- ions

  3. Looking at just sodium: 2Na -> 2Na+, so the ionic equation must be: 2Na -> 2Na+ + 2e-, meaning sodium has lost electrons & has been oxidised

  4. Looking at just chlorine: 2Cl- -> 2Cl-, meaning chlorine has not been oxidised or reduced

  5. Looking at just hydrogen: 2H+ -> H2, so the ionic equation must be: 2H+ + 2e- -> H2, meaning hydrogen has gained electrons so has been reduced
• Electrolysis

  The process of breaking down an ionic substance into its elements by passing an electric current through it
• 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 solution

  4. Substance is broken down into elements
• Electrolyte


  The substance being broken down during electrolysis
• What happens during electrolysis

  1. Positively charged ions move to the negative electrode (cathode)

  2. Negatively charged ions move to the positive electrode (anode)

  3. Ions are discharged at the electrodes producing elements
• Electrolysis of molten ionic compounds


  1. Metal is produced at the cathode

  2. Non-metal is produced at the anode
• This is because the metal is the positive ions and the non-metal is the negative ions
• Extracting metals by electrolysis


  • Metals more reactive than carbon are extracted by electrolysis of molten compounds
  • Large amounts of energy are used to melt the compounds and produce the electrical current
• Aluminium extraction by electrolysis

  • Molten mixture of aluminium oxide and cryolite is electrolysed using carbon as the positive electrode
  • Oxygen reacts with the carbon electrodes, forming carbon dioxide and causing them to gradually burn away
• 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. Hydrogen is produced at the cathode unless the metal is less reactive than hydrogen

  3. One of the halide ions is produced at the anode if present, otherwise oxygen is formed
• This happens because water molecules break down in the aqueous solution producing H+ and OH- ions that are discharged
• Half equations


  Represent the reactions at the electrodes, with the small number always the same as the 2 larger numbers, and electrons represented as 'e-'
• Writing half equations for the reactions at each electrode


  1. Negative electrode: X+ + e- -> X (positive ions are reduced)

  2. Positive electrode: X- -> e- + X (negative ions are oxidised)
• Acid + metal


  Produces salt + hydrogen
• Reactions of acids with metals

  • They are redox reactions - one substance is reduced and another is oxidised
• Identifying oxidation and reduction in acid-metal reactions

  1. Look at electrons gained and lost (following OIL RIG)

  2. Mg: Mg -> Mg2+ + 2e- (oxidised)

  3. H+: 2H+ + 2e- -> H2 (reduced)
• Neutralisation


  Acids are neutralised by alkalis (soluble metal hydroxides) and bases (insoluble metal hydroxides and metal oxides) to produce salts and water
• Acids + metal carbonates

  Produce salt + water + carbon dioxide
• Salts produced from acid reactions


  • Chlorides (from HCl)
  • Nitrates (from HNO3)
  • Sulfates (from H2SO4)
• Charges on ions in salts

  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 to acid until excess sinks, filter, evaporate to crystallise
• pH scale

  Measures acidity/alkalinity, pH 7 is neutral, pH<7 is acidic, pH>7 is alkaline
• Neutralisation reaction


  H+(aq) + OH-(aq) -> H2O(l)
• Titration


  1. Measure volumes of acid and alkali that react using a burette and indicator

  2. Repeat to get precise results
• Titration calculations


  • 1 dm3 = 1000 cm3
  • Moles = volume x concentration