Electrochemistry

1. Using and Controlling Reactions Electrochemistry

2. Redox Half Equations • Assign oxidation numbers and balance atom whose oxidation number changes • Balance oxygen by adding water • Balance hydrogen by adding H+ • Balance charges by adding electrons (always on the same side as the added H+) • Check the equation

3. Balancing Redox Equations • Multiply one or both equations by appropriate numbers so that the number of electrons lost or gained in each equation is equal • Add the two equations cancelling electrons (and other species as necessary) • CHECK THE EQUATION!!!!!!!!!!!!!!

4. Electrochemical cells Galvanic cells Electrolytic cells Primary cells Secondary cells Fuel cells Electrochemistry

5. Galvanic Cells • Produce electrical energy from spontaneous redox reactions • Consist of two half cells (metal or solution) where the oxidising agent and reducing agent are not in contact with each other. • The two half cells are connected via a conducting wire (connects the electrodes) and the salt bridge (connects the solutions)

6. Galvanic Cells • Salt bridge consists of a concentrated solution of a salt which is not easily oxidised or reduced • Oxidation occurs at the ANODE (negative electrode) • Reduction occurs at the CATHODE (positive electrode)

7. Galvanic Cells • Electrons flow from anode to cathode through the external wire • Positive ions move from the salt bridge into the reduction half cell • Negative ions move from the salt bridge into the oxidation half cell

8. Metal Half Cells • Solid metal electrode • Solution containing ions of the same metal (usually a sulfate salt) • More reactive metal is oxidised at the anode: M  Mx+ + xe • Less reactive metal is reduced at the cathode: Ny+ + ye  N • (x and y represent number of electrons gained or lost by metal/ metal ion)

9. Galvanic Cell using Metal Half Cells

10. Solution Half Cells • Inert electrodes (Graphite or Platinum) • The reacting solutions may contain an oxidant (e.g. MnO4–) or a reductant (e.g. I–) • Sulfuric acid is used to acidify solutions in half cell where necessary for a reaction to occur • Electrons are donated or accepted from the solution, not the electrode

11. Fuel Cells • Gaseous fuel (most often H2 gas) is oxidised at the anode. H2(g) 2H+(aq) + 2e • Oxidant (oxygen gas) is reduced at the cathode. O2(g) + 4H+(aq) + 4e  2H2O(l) • Overall reaction 2H2(g) + O2(g) 2H2O(l)

12. Fuel Cells • Electrodes: Porous graphite, containing platinum based catalyst. (To increase rate of reaction) • Salt Bridge: Five main types which identifies the fuel cell type. (Alkaline, Solid Polymer (PEM), Phosphoric acid, Molten carbonate, Solid oxide) These allow passage of ions but block the passage of electrons.

13. Advantages of Fuel Cells • High operating efficiency • Environmentally friendly (don’t produce SO2, NOx) • Quiet and reliable. Will run as long as the fuel is available and require minimal maintenance. • Better mass to power output compared to conventional galvanic cells • Fuel and oxidant readily available

14. Advantages of Fuel Cells • Products are removed as formed, rather than staying inside the cell. • Require minimal maintenance as there are no moving parts. • Can be used for a large range of applications.

15. Disadvantages • High purity fuels and oxidants are expensive and are often produced using natural gas as a feedstock. • Impurities in the fuel can “poison” the catalyst in the electrodes • Electrodes are expensive due to the catalyst • Many of the electrolytes are corrosive • Rate of reaction is slow. Medium to high temperatures are required for the cell to function. • Safety and Storage of Hydrogen?

16. Mercedes NECAR Hydrogen Fuel Cell Car http://www.cardesignonline.com/technology/necar-fuel-cell.php

17. Hydrogen Fuel Cell Bicycles http://www.alternative-energy-news.info/hydrogen-fuel-cell-bikes

18. Portable fuel cell powered by water and Aluminium http://pinktentacle.com/2006/04/portable-fuel-cell-powered-by-water-and-aluminum/

19. Sony Exhibiting Hybrid Fuel Cell Batteries in Tokyo http://cleantechnica.com/2009/02/26/sony-exhibiting-hybrid-fuel-cell-batteries-in-tokyo/

20. World's smallest fuel cell promises greener gadgets http://www.newscientist.com/article/dn16370-worlds-smallest-fuel-cell-promises-greener-gadgets.html

21. Similarities between Fuel Cells and Conventional Cells • Redox reactions used to produce direct current. • Electrolyte between electrodes. • No pollutants emitted. • Anode is negative and cathode is positive electrode.

22. Differences between Fuel Cells and conventional cells

23. Rechargeable cells • Referred to as storage cells or accumulators • Act as galvanic cells when discharging • During recharging an electric current reforms the original substances • Common types include the lead acid accumulator and the NICAD (nickel cadmium cell)

24. Example: Lead Acid Accumulator • Power source in motor vehicles • Six lead acid cells connected in series (generate 2V each) • Anode: Lead • Cathode: Lead oxide on lead • Electrolyte: Sulfuric acid (38%w/v)

25. Lead Acid Accumulator • Discharging • Anode(-): Pb(s) Pb2+ + 2e • Cathode(+): PbO2(s)+ 4H+(aq)+ 2e  Pb2+(aq)+ 2H2O(l) • The lead ions react with sulfate ions to form insoluble lead sulfate: Pb2+(aq) + SO42-(aq) PbSO4(s)

26. Lead Acid Accumulator • Overall: PbO2(s)+ Pb(s)+ 2SO42-(aq)+ 4H+ 2PbSO4(s)+ 2H2O(l) • Anode, cathode and electrolyte are consumed in the reaction • The state of charge/discharge of the battery can be measured by the density of the electrolyte

27. Lead Acid Accumulator • Charging: • Anode(-) when discharging becomes the cathode(-) when charging: PbSO4(s) + 2e  Pb(s) + SO42-(aq) • Cathode(+) when discharging becomes the anode(+) when charging: PbSO4(s)+ 2H2O(l) PbO2(s)+ 4H+(aq)+ SO42-(aq)+2e

28. Lead Acid Accumulator • Overall:(opposite reaction to discharging) 2PbSO4(s)+2H2O(l) PbO2(s)+ Pb(s)+2SO42-(aq)+4H+ • This regenerates the anode and cathode and increases the density of the electrolyte

29. Electrolytic Cells • Change electrical energy into chemical energy • Cause a non spontaneous redox reaction to occur • Electrodes can be reactive or inert • Electrolyte is a solution or molten liquid. The chemicals reactivity related to the reactivity of water determines which is used.

31. Electrolytic Cells • Oxidation occurs at the anode (+) and reduction occurs at the cathode (-) • If the electrolyte is molten then the anions (-ve ion) are oxidised at the anode and the cations (+ve ion) are reduced at the cathode. • If the electrolyte is aqueous then the reactions could involve the cations, anions or water.

32. Electrolytic Cells • Reduction • Water will be reduced in preference to the metals in the activity series Al and above: • 2H2O + 2e → 2OH- + H2 • Zn and below will undergo reduction in an aqueous solution: • M2+ + 2e → M (M represents metal)

33. Electrolytic Cells • Oxidation • Chloride, bromide and iodide are oxidised in preference to water: • 2X- → X2 + 2e (X represents halogen) • Nitrate and sulfate ions will not oxidise. (N and S already in max oxidation state) • When these ions are present water will oxidise: • H2O → 4H+ + O2 + 4e

34. Uses of Electrolytic Cells • Extraction of metals from molten salts • Refining metals • Electroplating for protection or decoration • Recharging secondary cells • Production of chemicals (NaOH, H2, Cl2,O2)