1. Using and Controlling Reactions 1

2. 1. Assign oxidation numbers and balance atom whose oxidation number changes 2. Balance oxygen by adding water 3. Balance hydrogen by adding H + 4. Balance charges by adding electrons (always on the same side as the added H + ) 5. Check the equation 2

3. 1. Multiply one or both equations by appropriate numbers so that the number of electrons lost or gained in each equation is equal 2. Add the two equations cancelling electrons (and other species as necessary) 3. CHECK THE EQUATION!!!!!!!!!!!!!! 3

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

5.  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) 5

6.  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) 6

7.  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 7

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

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10.  Inert electrodes (Graphite or Platinum)  The reacting solutions may contain an oxidant (e.g. MnO 4 – ) 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 10

11.  Gaseous fuel (most often H 2 gas) is oxidised at the anode. H 2(g)  2H + (aq) + 2e  Oxidant (oxygen gas) is reduced at the cathode. O 2(g) + 4H + (aq) + 4e  2H 2 O (l)  Overall reaction 2H 2(g) + O 2(g)  2H 2 O (l) 11

12.  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. 12

13.  High operating efficiency  Environmentally friendly (don’t produce SO 2, NO x )  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 13

14.  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. 14

15.  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? 15

16. 16 http://www.cardesignonline.com/technology/necar-fuel-cell.php

17. 17 http://www.alternative-energy-news.info/hydrogen-fuel-cell-bikes

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

19. 19 http://cleantechnica.com/2009/02/26/sony-exhibiting-hybrid-fuel-cell-batteries-in- tokyo/

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

21.  Redox reactions used to produce direct current.  Electrolyte between electrodes.  No pollutants emitted.  Anode is negative and cathode is positive electrode. 21

22. Conventional galvanic cells Fuel cells Limited quantities of reactants stored in cell Continuous external supply of reactants Must be discarded or recharged when fully discharged Never discharge or run down Limited life Virtually unlimited life 22

23.  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) 23

24.  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) 24

25.  Discharging  Anode(-): Pb (s)  Pb 2+ + 2e  Cathode(+): PbO 2(s) + 4H + (aq) + 2e  Pb 2+ (aq) + 2H 2 O (l)  The lead ions react with sulfate ions to form insoluble lead sulfate: Pb 2+ (aq) + SO 4 2- (aq)  PbSO 4(s) 25

26.  Overall: PbO 2(s) + Pb (s) + 2SO 4 2- (aq) + 4H +  2PbSO 4(s) + 2H 2 O (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 26

27.  Charging:  Anode(-) when discharging becomes the cathode(-) when charging: PbSO 4(s) + 2e  Pb (s) + SO 4 2- (aq)  Cathode(+) when discharging becomes the anode(+) when charging: PbSO 4(s) + 2H 2 O (l)  PbO 2(s) + 4H + (aq) + SO 4 2- (aq) +2e 27

28.  Overall: (opposite reaction to discharging) 2PbSO 4(s) +2H 2 O (l)  PbO 2(s) + Pb (s) +2SO 4 2- (aq) +4H +  This regenerates the anode and cathode and increases the density of the electrolyte 28

29.  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. 29

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31.  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. 31

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

33.  Oxidation  Chloride, bromide and iodide are oxidised in preference to water:  2X - → X 2 + 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:  H 2 O → 4H + + O 2 + 4e 33

34.  Extraction of metals from molten salts  Refining metals  Electroplating for protection or decoration  Recharging secondary cells  Production of chemicals (NaOH, H 2, Cl 2, O 2 ) 34

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