Oxidation and Reduction

1. Oxidation and Reduction Chemistry 3B 2014

2. Oxidation and Reduction • apply the table of Standard Reductions Potentials to determine the relative strength of oxidising and reducing agents to predict reaction tendency • apply oxidation numbers to identify redox equations and/or oxidants and reductants • identify by name and/or formula common oxidising and reducing agents including O2, Cl2, MnO4–, Cr2O72–, ClO–, H+, concentrated sulfuric acid, concentrated nitric acid and common reducing agents (reductants) including Zn, C, H2, Fe2+ , C2O42– • write and balance oxidation-reduction half equations in acidic conditions • write balanced oxidation-reduction equations • describe and explain the role of the following in the operation of an electrochemical (galvanic) cell: anode processes cathode processes electrolyte salt bridge and ion migration electron flow in external circuit • describe the electrical potential of a galvanic cell as the ability of a cell to produce an electric current • describe and explain how an electrochemical cell can be considered as two half-cells • describe the role of the hydrogen half-cell in the table of Standard Reduction Potentials • describe the limitations of Standard Reduction Potentials table.

3. Assigning Oxidation Numbers

4. Assigning Oxidation Numbers • The oxidation number of a Group IA element in a compound is +1. • The oxidation number of a Group IIA element in a compound is +2. • The sum of the oxidation numbers of all of the atoms in a neutral compound is 0. For example SO2 = 0 • The sum of the oxidation numbers in a polyatomic ion is equal to the charge of the ion. For example SO42- = 2-

5. Redox Revisited • The chemical changes that occur when electronsare transferred between reactants are called oxidation – reduction reactions. • Oxidation reactions are always accompanied by a reduction reaction. • The substance that donates electrons in a redox reaction is the REDUCING AGENT. • The substance that takes electrons in a redox reaction is the OXIDISING AGENT.

6. Identifying redox reactions Reduced  Oxidising agent (Cl2) Cl2 + 2Na  2NaCl (0) (0) (+1)(-1) Oxidised Reducing agent (Na) The oxidised species is identified by an increase in its oxidation number while the reduced species shows a decrease in its oxidation number.

7. Common oxidising and reducing agents

8. Increasing strength of reducing agents Increasing strength of oxidising agents

9. Half equations • An oxidation half equation shows the species which is oxidised (lost electrons). • It has electrons on the products side of the half equation. 2Hg  Hg22+ + 2e- • The reduction half equation shows the species which is reduced (gained electrons). • It has the electrons on the reactant side of the half equation. Cl2 + 2e- 2Cl- A useful mnemonic to remember is OIL RIG

10. Half equations • Together the two half equations describe the overall redox reaction. 2Hg  Hg22+ + 2e- Cl2 + 2e- 2Cl- 2Hg + Cl2 Hg22+ + 2Cl-

11. Balancing half equations in acidic conditions (1) Balance the following redox reaction in acidic conditions. Cr2O7−2(aq)+HNO2(aq)→ Cr3+(aq) + NO3− (aq) Step 1: Separate the half-reactions. Cr2O7−2(aq)→Cr3+(aq) HNO2(aq)→NO3−(aq) Step 2: Balance elements other than O and H. Cr2O7−2(aq)→2Cr3+(aq) HNO2(aq)→NO3− (aq)

12. Balancing half equations in acidic conditions (2) Step 3: Add H2O to balance oxygen. Cr2O7−2(aq)→2Cr3+(aq)+7H2O(l) HNO2(aq)+H2O(l)→NO3−(aq) Step 4: Balance hydrogen by adding H+. 14H+(aq)+Cr2O7−2 (aq)→2Cr3+(aq)+7H2O(l) HNO2(aq)+H2O(l)→3H+(aq)+NO3−(aq)

13. Balancing half equations in acidic conditions (3) Step 5: Balance the charge of each equation with electrons. HNO2(aq)+H2O(l)→3H+(aq)+NO3-(aq)+2e− 6e−+14H+(aq)+Cr2O7−2(aq)→2Cr3+(aq)+7H2O(l). Step 6: Scale the reactions so that the electrons are equal. 3∗[HNO2(aq) + H2O(l)→3H+(aq) +NO3−(aq) +2e−] ⇒ 3HNO2(aq) +3H2O(l)→9H+(aq)+3NO3− (aq)+6e−

14. Balancing half equations in acidic conditions (4) Step 7: Add the reactions and cancel out common terms. 3HNO2(aq)+3H2O(l)→9H+(aq)+3NO3−(aq)+6e− 6e− +14H+(aq)+Cr2O7−2(aq)→2Cr3+(aq)+7H2O(l) 3HNO2(aq)+3H2O(l)+6e−+14H+(aq)+Cr2O7−2 (aq)→9H+(aq)+3NO3− (aq)+6e−+2Cr3+(aq)+7H2O(l) The electrons cancel out as well as 3 water molecules and 9 H+. This leaves the balanced net reaction of: 3HNO2(aq) +5H+(aq)+ Cr2O7−2(aq)→ 3NO3−(aq)+2Cr3+ (aq)+4H2O(l)

15. Electrochemical Cells • Galvanic cells, voltaic cells and electrochemical cells are commonly referred to as batteries. Although technically a battery is defined as two or more electric cells connected in series to produce a steady flow of current. • They use a redox reaction to produce a voltage (potential difference or EMF) that results in an electric current.

16. Electrochemical Cells • Electrochemical cells are composed of two electrodes(solid electrical conductors) and at least one electrolyte(aqueous electrical conductor). The electrolyte is often a moist paste (just enough water is added so that the ions can move). Sometimes one electrode is the cell container. • The positive electrode is the cathodeand the negative electrode is the anode. The electrons flow through the external circuit from the anode to the cathode. • A porous boundary or salt bridge separates the two electrolytes while still allowing ions to flow to maintain cell neutrality.

17. Daniell Cell The Daniell cell uses copper sulphate (Cu2+) as the oxidising agent and zinc metal (Zn) as the reducing agent. The zinc electrode becomes oxidised releasing zinc ions and electrons. These electrons flow from zinc electrode, through the external conducting path, to the copper electrode. Here copper ions in the electrolyte move to the copper electrode where they gain electrons from the oxidation of zinc, to be reduced to metallic copper. A salt bridge completes the circuit allowing ions to move between the cathode and anode.

18. Voltaic Virtual Cell Lab

19. Standard Reduction Potentials • The standard reduction potential is the potential in volts generated by a reduction half-reaction compared to the standard hydrogen electrode at 25 °C (298 K) , 100.0 kPa and a concentration of 1.0 molL-1. • Standard reduction potentials are denoted by the variable E0(E nought). • All standard reduction potentials are relative to the standard hydrogen half cell being 0.00V. • This means that all standard reduction potentials that are positive are stronger oxidizing agents than hydrogen ions and all standard reduction potentials that are negative are weaker.

20. Standard Reduction Potential Table • The organisation of the Standard Reduction Potential(SRP) Table allows it to be used to predict if a given redox reaction will be spontaneous. • If the oxidising agent is situated above the reducing agent then the reaction will be spontaneous. • To calculate the Eo (cell) the Eo values for the oxidation and reduction half reaction must be added. • A positive value means the redox reaction is spontaneous, a negative value means the reaction will not occur.

21. Example Is the following redox reaction spontaneous? 2Al + 6H+ 3H2 + 2Al3+ 2Al  2Al3+ + 6e- +1.66 6H+ + 6e-  3H2 0.00 2Al + 6H+ 3H2 + 2Al3+ +1.66 Yes the reaction is spontaneous

22. Redox in Action – Primary Cell • Primary batteries are non-rechargeable and disposable and are commonly used by smoke detectors, flashlights, and most remote controls. • The dry cell produces a maximum voltage of 1.5V, however this decreases over time as reactant concentration decreases. They also have a low energy to mass ratio or energy density. • The cell has negligible environmental impact and cells are disposed of in household rubbish bins.

23. Primary Cell Example - Dry Cell The overall cell reactions for a dry cell are: Anode: Zn(s) -> Zn2+(aq) + 2e- Cathode: 2NH4+(aq) + 2MnO2(s) + 2e- -> Mn2O3(s) + H2O(l) + 2NH3(aq)

24. Redox in Action – Secondary Cells • Secondary batteries are rechargeable. These batteries undergo electrochemical reactions that can be readily reversed. Secondary batteries are used in car batteries and portable electronic devices. • The lead-acid cell produces a voltage of 2V. In a car battery six of these cells are connected in series to produce a voltage of approximately 12V. • The cell is recharged by applying a direct current of slightly greater than 2V to each cell.

25. Example of a Secondary Cell - Lead acid cell The anode grid contains spongy lead, while the cathode grid is packed with powdered lead dioxide. The powdered nature of both these reagents allows for a fast reaction rate because of their high surface area. This allows the cell to produce high currents. The electrolyte is concentrated sulphuric acid. Anode: Pb+ SO42-(aq)  → PbSO4(s) + 2e-  Cathode: 4H+ + PbO2  + SO42-  + 2e- → PbSO4(s) + 2H2O The overall equation is Pb(s) +  PbO2(s) + 4H+ + 2SO4 2-(aq) → 2PbSO4(s) + 2H2O(l)   

26. Past Examination Question - 2013

27. Past Examination Question - 2013

28. Past Examination Question - 2013

29. Past Examination Question - 2013

30. Past Examination Question - 2013

31. Past Examination Question - 2013

32. Redox in action – Fuel Cells Fuel cells do not store the oxidising and reducing agents, instead these reactants are constantly fed into the cell to generate electricity. Three Mercedes-Benz model O530BZ “Citaro” hydrogen fuel-cell buses were used in Perth between 2004 and 2007 as part of an international trial to study the operation of alternative energy driven buses.

33. Example of a Fuel Cell – Alkaline hydrogen oxygen Anode: H2(g) + 2OH-(aq)  2H2O(l) + 2e- Cathode: O2(g) + 2H2O(l) + 4e-  4OH-(aq) Overall redox reaction: 2H2(g) + O2(g)  2H2O(l) The reducing agent H2 diffuses into the nickel electrode and is oxidised forming H+ ions and free electrons. Oxygen absorbed into the cathode then gains electrons forming OH- ions.

34. Redox in Action - corrosion • Metal corrosion is an electrochemical process where a metal is oxidised in the presence of agents such as water and air (O2). • Corrosion of iron (or rusting) involves the oxidation of Fe by air (O2) in a moist environment.

35. Corrosion of iron Considering the sketch of a water droplet, the oxidizing iron supplies electrons at the edge of the droplet to reduce oxygen from the air. The iron surface inside the droplet acts as the anode for the process Fe(s) -> + Fe2+(aq) + 2e- The electrons can move through the metallic iron to the outside of the droplet where O2(g) + 2H2O(l) + 4e- -> 4OH-(aq) Within the droplet, the hydroxide ions can move inward to react with the iron(II) ions moving from the oxidation region. Iron(II) hydroxide is precipitated. Fe2+(aq) + 2OH-(aq) -> Fe(OH)2(s) Rust is then quickly produced by the oxidation of the precipitate. 4Fe(OH)2(s) + O2(g) -> 2Fe2O3 •H2O(s) + 2H2O(l)

36. Factors affecting the rate of corrosion • Oxygen – higher concentration the faster the rate of corrosion • Water – absence of water corrosion will not occur • pH – the more acidic the environment the faster the rate of corrosion • Electrolytes – corrosion occurs at a faster rate in salty environments • Less reactive metals – contact with a more reactive metal prevents corrosion • Temperature- a higher temperature the faster the rate of corrosion

37. Sacrificial Anodes Sacrificial anodes are usually made from magnesium, zinc or aluminium; magnesium is used on-shore and in freshwater, and zinc and aluminium are used in salt water where resistance is lower.

38. Prevention of corrosion of iron. • Inert non-metallic coatings – grease, paint, plastic • Inert metallic coatings – with less reactive metal – tin or copper • Galvanising- coating iron with more reactive zinc which provides an oxide coating • Cathodic protection using a sacrificial anode – as per previous slide • Cathodic protection using a DC current-

39. Cathodic Protection • This involves making iron the cathode in the cell. • The cell consists of a DC (direct current) power source, an anode of scrap metal and the steel cathode. • The applied voltage makes the iron object negatively charged and so prevents oxidation. • The anode is replaced regularly.