4.4 CHEMISTRY

• 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. Magnesium: Mg -> Mg2+ + 2e-, Mg has lost electrons so Mg has been oxidised

  3. Hydrogen: 2H+ + 2e- -> H2, H has gained electrons, so H has been 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 salts, water and carbon dioxide
• Salts produced from acid reactions

  • Chlorides (XCl) from hydrochloric acid (HCl)
  • Nitrates (XNO3) from nitric acid (HNO3)
  • Sulfates (XSO4) from sulfuric acid (H2SO4)
• Charges on ions

  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

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

  2. Keep adding until excess solid sinks to bottom, then filter out excess solid

  3. Evaporate some water, then leave to evaporate slowly to produce crystals
• pH scale

  Measures acidity or alkalinity of a solution, 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 solutions that react with each other

  2. Wash burette, fill with acid, add alkali to flask, add indicator, add acid from burette until end-point reached

  3. Titre is difference between first and second burette readings

  4. Repeat for more precise results
• Titration calculations

  • 1 dm3 = 1000 cm3
  • One mole of a substance in grams is the same as its relative atomic mass in grams
• Calculating acid concentration from titration

  1. Convert volumes to dm3

  2. Calculate moles of NaOH from volume and concentration

  3. Use mole ratio from equation to calculate moles of HCl

  4. Calculate concentration of HCl from moles and volume
• Strong acid

  Completely ionised in aqueous solution, e.g. hydrochloric, nitric and sulfuric acids
• Weak acid

  Partially ionised in aqueous solution, e.g. ethanoic, citric and carbonic acids
• Acid strength

  Stronger an acid, lower the pH (for a given concentration of aqueous solutions)
• pH decreases by 1 unit

  H+ concentration of solution 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
• Electrolysis

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

  1. Ions are free to move in a molten or dissolved ionic substance

  2. Passing a current through the substance breaks it down into elements

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

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

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

  • The metal is produced at the cathode
  • The non-metal is produced at the anode
• Metals extracted by electrolysis

  Metals more reactive than carbon, which are too reactive to be extracted by reduction with carbon
• Extracting metals by electrolysis

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

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

  3. The positive electrodes need to be continually replaced as oxygen reacts with the carbon, forming carbon dioxide
• Electrolysis of aqueous solutions

  • The ions discharged depend on the relative reactivity of the elements
  • At the cathode, hydrogen is produced unless the metal is less reactive than hydrogen
  • At the anode, if OH- and halide ions are present, one of the halide ions will be produced, otherwise oxygen is formed

  
• Half equations

  • Represent the reactions at the electrodes
  • The small number is always the same as the 2 larger numbers within the equation
  • Electrons are represented by the symbol '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
• 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
• Reactivity series of metals

  • Potassium
  • Sodium
  • Lithium
  • Calcium
  • Magnesium
  • Aluminium
  • Carbon
  • Zinc
  • Iron
  • Hydrogen
  • Copper
  • Silver
  • Gold
• Reactivity series of metals based on reaction with water


  • Potassium (violent)
  • Sodium (very quick)
  • Lithium (quick)
  • Calcium (more slow)
• Reactivity series of metals based on reaction 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
• The charges on each side of an ionic equation should add up to the same number