Mr Cole Chemistry

Catalysts Questions

  • Questions
  • Answers
  1. State whether the following reactions are homogenous or heterogenous reactions.
    1. The catalytic cracking of C20H42 using zeolite.
    2. The oxidation of sulfur dioxide to sulfur trioxide using vanadium (V) oxide.
    3. The reaction between Iodide and peroxodisulfate, catalysed by Fe2+
    4. The Haber process
  2. Define the following terms:
    1. Homogenous
    2. heterogenous
  3. Vanadium (V) oxide is used as a catalyst in the Contact Process, where it is used to catalyse the oxidation of sulfur dioxide to sulfur trioxide.
    1. Write an equation for the oxidation of sulfur dioxide to sulfur trioxide.
    2. Write two equations showing how vanadium (V) oxide behaves as a catalyst.
    3. State how you can tell from the reaction that the vanadium (V) oxide is a catalyst.
    4. Describe the process of catalysis in this reaction.
    5. The presence of lead in the sulfur raw material can cause the reaction to slow down. Explain why this happens.
    6. Sketch a reaction profile of the reaction with and without the catalyst. The reaction is exothermic.
  4. C2O42- is oxidised by  MnO4– in an autocatalyzed reaction.
    1. Write an equation for oxidation of C2O42- by MnO4–Define the term autocatalysis.
    2. Explain why the reaction is slow at the start of the reaction.
    3. Explain why the reaction speeds up after some time. Include equations in your answer.
    4. Sketch a graph showing the change in Mn2+ concentration over time.
  5. Transition metals and their ions make good catalysts.
    1. State the feature of transition metals that makes them good catalysts.
    2. Observe the graph below and use it to explain why rhodium and palladium are the best catalysts.
  6. Peroxodisulfate ions are reduced to sulfate ions by iodide ions in a reaction that is catalysed by iron (II) ions.
    1. Write the half equations for:
      1. Oxidation of iodide ions.
      2. Reduction of S2O82- ions.
    2. Write the ionic equation for the reaction.
    3. Write two equations showing how Fe2+ behaves as a catalyst for the reaction.
    4. State two reasons why iron (II) is a good choice for a catalyst for the reaction.
    5. Using the electrode potentials provided, explain how Fe2+ catalyses the reaction.
      S2O82– + 2e– ⇌ 2SO42– E° = +2.01 V
      Fe3+ + e– ⇌ Fe2+ E° = +0.77
      I2 + 2e– ⇌ 2I– E° = +0.54 V
  7. The Haber Process is catalysed by a transition metal.
    1. State the identity of the metal that catalyses the reaction.
    2. Explain why the catalyst does not affect the yield of ammonia.
    3. Explain how the metal catalyses the reaction.
    4. State why the reactants should be purified before they come into contact with the metal catalyst.
  8. Sulfur dioxide is removed from flue gases by reacting with a neutralising agent. The reaction produces calcium sulfate (IV).
    1. Write an equation for the reaction.
    2. The sulfate (IV) ions are oxidised by oxygen in the air to produce sulfate (VI). Write an equation for the reaction.
    3. An alternative mechanism involves the use of Co3+ ions to catalyse the reaction.
      1. Write an equation for the oxidation of SO32- to SO42- using acidified redox.
      2. Write an equation for the reduction of acidified Co3+ ions using oxygen.
      3. Explain how Co3+ behaves as a catalyst in this reaction.
  1. State whether the following reactions are homogenous or heterogenous reactions.
    1. The catalytic cracking of C20H42 using zeolite.
      Heterogenous
    2. The oxidation of sulfur dioxide to sulfur trioxide using vanadium (V) oxide.
      Heterogenous
    3. The reaction between Iodide and peroxodisulfate, catalysed by Fe2+
      Homogenous
    4. The Haber process
      Heterogenous
  2. Define the following terms:
    1. Homogenous
      The catalyst is in the same phase as the reactants and products
    2. heterogenous
      The catalyst is in a different phase to the reactants and products
  3. Vanadium (V) oxide is used as a catalyst in the Contact Process, where it is used to catalyse the oxidation of sulfur dioxide to sulfur trioxide.
    1. Write an equation for the oxidation of sulfur dioxide to sulfur trioxide.
      SO2 + ½ O2 –> SO3
    2. Write two equations showing how vanadium (V) oxide behaves as a catalyst.
      SO2 + V2O5 –> SO3 + V2O4
      V2O4 + ½O2 –> V2O5
    3. State how you can tell from the reaction that the vanadium (V) oxide is a catalyst.
      V2O5 reacts in the first step but then regenerates in the second step.
    4. Describe the process of catalysis in this reaction.
      Step 1: Sulfur dioxide and oxygen adsorb to the active site on the surface of the vanadium (V) oxide.
      Step 2: Sulfur dioxide is oxidised to sulfur trioxide by the vanadium (V) oxide.
      Step 3: Sulfur trioxide desorbs from the active site on the surface of the catalyst.
      Step 4: Oxygen reacts with the vanadium (IV) oxide, oxidising it back to vanadium (V) oxide.
      Steps 3 and 4 can happen either way round.
    5. The presence of lead in the sulfur raw material can cause the reaction to slow down. Explain why this happens.
      Lead in the sulfur raw material can end up blocking the active site on the catalyst, poisoning the catalyst and reducing the rate of reaction.
    6. Sketch a reaction profile of the reaction with and without the catalyst. The reaction is exothermic.
  4. C2O42- is oxidised by  MnO4– in an autocatalyzed reaction.
    1. Write an equation for oxidation of C2O42- by MnO4–Define the term autocatalysis.
      (C2O42- –> 2CO2 + 2e–)
      (MnO4– 8H+ + 5e– –> Mn2+ + 4H2O)
      2MnO4– 16H+ + 5C2O42- –> 2Mn2+ + 8H2O + 10CO2
    2. Explain why the reaction is slow at the start of the reaction.
      The reaction is catalysed by one of the products from the reaction.
    3. Explain why the reaction speeds up after some time. Include equations in your answer.
      C2O42- ions and MnO4– ions repel each other due to both being negative.
    4. Sketch a graph showing the change in Mn2+ concentration over time.
      Mn2+ catalyses the reaction. Mn2+ is positive and therefore is attracted to the negative MnO4- ions. It reduces the MnO4- to form Mn3+ which then oxidises the C2O42-. The Mn2+ is regenerated in the second step.
      (MnO4- 8H+ + 4e- –> Mn3+ + 4H2O)
      (Mn2+ –> Mn3+ + e-)
      MnO4- 8H+ + 4Mn2+ –> 5Mn3+ + 4H2O
      (C2O42- –> CO2 + 2e-)
      (Mn3+ –> Mn2+ + e-)
      C2O42- + Mn3+ –> CO2 + 2Mn2+
  5. Transition metals and their ions make good catalysts.
    1. State the feature of transition metals that makes them good catalysts.
      Variable oxidation states
    2. Observe the graph below and use it to explain why rhodium and palladium are the best catalysts.
      Their adsorption energy is neither too low nor too high. It is low enough that the reactants can adsorb to the surface but low enough that the products can desorb.
  6. Peroxodisulfate ions are reduced to sulfate ions by iodide ions in a reaction that is catalysed by iron (II) ions.
    1. Write the half equations for:
      1. Oxidation of iodide ions.
        2I– –> I2 + 2e–
      2. Reduction of S2O82- ions.
        S2O82- + 2e– –> 2SO42-
    2. Write the ionic equation for the reaction.
      S2O82- + 2I– –> 2SO42- + I2
    3. Write two equations showing how Fe2+ behaves as a catalyst for the reaction.
      S2O82- + 2Fe2+ –> 2SO42- + Fe3+
      2Fe3+ + 2I– –> 2Fe2+ + I2
    4. State two reasons why iron (II) is a good choice for a catalyst for the reaction.
      It has variable oxidation states and is a positive ion (and is therefore attracted to the negative peroxodisulfate ions)
    5. Using the electrode potentials provided, explain how Fe2+ catalyses the reaction.
      S2O82– + 2e– ⇌ 2SO42– E° = +2.01 V
      Fe3+ + e– ⇌ Fe2+ E° = +0.77
      I2 + 2e– ⇌ 2I– E° = +0.54 V
      The Eo for S2O82-/SO42- is higher than the Eo for Fe3+/Fe2+ so S2O82- will be reduced and Fe2+ will be oxidised.
      The Eo for I2/I– is 0.54, so Fe3+ will be reduced and I– oxidised.
  7. The Haber Process is catalysed by a transition metal.
    1. State the identity of the metal that catalyses the reaction.
      Fe
    2. Explain why the catalyst does not affect the yield of ammonia.
      Catalysts speed up the rate of the forward and backwards reactions equally.
    3. Explain how the metal catalyses the reaction.
      N2 and H2 adsorb onto the active site on the surface of the iron particle.
      Bonds break in the reactants, bonds form between the atoms to form the products.
      Products desorb from the active site.
    4. State why the reactants should be purified before they come into contact with the metal catalyst.
      Impurities can poison the active site, blocking it and preventing catalysis.
  8. Sulfur dioxide is removed from flue gases by reacting with a neutralising agent. The reaction produces calcium sulfate (IV).
    1. Write an equation for the reaction.
      SO2 + CaCO3 –> CaSO3 + CO2
      OR
      SO2 + CaO –> CaSO3
    2. The sulfate (IV) ions are oxidised by oxygen in the air to produce sulfate (VI). Write an equation for the reaction.
      CaSO3 + ½ O2 –> CaSO4
      OR
      SO32- + ½ O2 –> SO42-
    3. An alternative mechanism involves the use of Co3+ ions to catalyse the reaction.
      1. Write an equation for the oxidation of SO32- to SO42- using acidified redox.
        (SO32- + H2O –> SO42- + 2H+ + 2e–)
        (Co3+ + e– –> Co2+)
        SO32- + H2O +         2Co3+ –> SO42- + 2H+ + Co2+
      2. Write an equation for the reduction of acidified Co3+ ions using oxygen.
        Co2+ + ½O2 + 2H+ –> Co3+ + H2O
      3. Explain how Co3+ behaves as a catalyst in this reaction.
        eCo3+ oxidises the SO32-, getting reduced in the process. The Co3+ is regenerated in the second step by reacting with oxygen.ther
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