Redox Reactions

Redox reactions, in which oxidation and reduction occur together, are behind processes such as combustion, respiration, photosynthesis, corrosion, metallurgy and the working of batteries. The chapter moves from the classical idea of oxidation as addition of oxygen or removal of hydrogen to the electronic concept, in which oxidation is the loss of electrons and reduction is the gain of electrons, as in Zn + Cu2+ -> Zn2+ + Cu. You will learn the rules for assigning oxidation numbers and the Stock notation, such as iron(III) oxide. Redox reactions are classified as combination, decomposition, displacement and disproportionation reactions. You will practise balancing redox equations by the oxidation number method and the half-reaction (ion-electron) method in acidic and basic media, for example the reaction of permanganate with iron(II) ions. The chapter explains redox titrations using potassium permanganate and potassium dichromate as self-indicating or externally indicated oxidants, and iodometric titrations. Finally, it introduces electrode processes: the Daniell cell, the standard hydrogen electrode, standard electrode potentials and the electrochemical series, which predict whether a metal can displace another from solution.

Detailed Worksheet: Redox Reactions Section A - Definitions (10 marks) 1. Define oxidation and reduction in terms of electron transfer. (2 marks) 2. What is the oxidation number? State two rules for assigning it. (2 marks) 3. What is a disproportionation reaction? Give an example. (2 marks) 4. What is Stock notation? Write the Stock names of FeO and Fe2O3. (2 marks) 5. What is the standard hydrogen electrode? (2 marks) Section B - Calculations and Applications (15 marks) 6. Calculate the oxidation number of the named element in each: Cr in K2Cr2O7, Mn in KMnO4, S in H2SO4, S in Na2S4O6, Fe in Fe3O4, N in NH4+. (3 marks) 7. Balance the following redox reaction in acidic medium by the ion-electron method: MnO4- + Fe2+ -> Mn2+ + Fe3+. State how many electrons are transferred per MnO4- ion. (3 marks) 8. Balance the following in acidic medium by the oxidation number method: Cr2O7 2- + SO3 2- -> Cr3+ + SO4 2-. Identify the oxidant and reductant. (3 marks) 9. 25.0 mL of 0.020 M KMnO4 is required to oxidise a solution of Fe2+ ions in acid. Using the balanced equation MnO4- + 5Fe2+ + 8H+ -> Mn2+ + 5Fe3+ + 4H2O, calculate the moles of Fe2+ oxidised and the mass of iron (Fe = 56). (3 marks) 10. The standard electrode potentials are Zn2+/Zn = -0.76 V and Cu2+/Cu = +0.34 V. Calculate the standard EMF of the Daniell cell and write the cell reaction. Which electrode is the anode? (3 marks) Section C - Diagrams (10 marks) 11. Draw a labelled diagram of the Daniell cell showing the electrodes, salt bridge and direction of electron flow. (4 marks) 12. Draw a labelled diagram of the standard hydrogen electrode. (3 marks) 13. Draw a flowchart showing the steps for balancing a redox reaction by the half-reaction method. (3 marks) Section D - Analysis and Higher-order Thinking (15 marks) 14. Explain the types of redox reactions with examples: combination, decomposition, metal and non-metal displacement, and disproportionation. Why can fluorine not show disproportionation? (5 marks) 15. Explain why KMnO4 acts as a self-indicator in redox titrations, while K2Cr2O7 needs an indicator. Describe the principle of iodometric titration. (5 marks) 16. Using standard electrode potentials, explain why zinc displaces copper from copper sulphate solution but copper cannot displace zinc. Arrange the metals Zn, Cu, Fe and Ag in order of their reducing power, given E zero values of -0.76 V, +0.34 V, -0.44 V and +0.80 V. (5 marks) Instructions: Time allowed 2 hours. Attempt all sections. Write balanced equations with charges. Show all calculation steps.