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ElectrochemistryUsing and Controlling Reactions
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Redox Half Equations
1. Assign oxidation numbers and balance atom whose oxidation number changes
2. Balance oxygen by adding water3. Balance hydrogen by adding H+ 4. Balance charges by adding
electrons (always on the same side as the added H+)
5. Check the equation
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Balancing Redox Equations
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!!!!!!!!!!!!!!
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Electrochemistry
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Electrochemicalcells
Galvaniccells
Electrolyticcells
Primarycells
Secondarycells
Fuelcells
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)
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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)
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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
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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)
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Galvanic Cell using Metal Half Cells
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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
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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 reaction2H2(g) + O2(g) 2H2O(l)
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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.
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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 available13
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.
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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? 15
Mercedes NECAR Hydrogen Fuel Cell Car
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http://www.cardesignonline.com/technology/necar-fuel-cell.php
Hydrogen Fuel Cell Bicycles
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http://www.alternative-energy-news.info/hydrogen-fuel-cell-bikes
Portable fuel cell powered by water and Aluminium
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http://pinktentacle.com/2006/04/portable-fuel-cell-powered-by-water-and-aluminum/
Sony Exhibiting Hybrid Fuel Cell Batteries in Tokyo
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http://cleantechnica.com/2009/02/26/sony-exhibiting-hybrid-fuel-cell-batteries-in-tokyo/
World's smallest fuel cell promises greener gadgets
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http://www.newscientist.com/article/dn16370-worlds-smallest-fuel-cell-promises-greener-gadgets.html
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.
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Differences between Fuel Cells and conventional cells
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
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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)
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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)
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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) + SO4
2-(aq) PbSO4(s) 25
Lead Acid Accumulator
Overall:PbO2(s)+ Pb(s)+ 2SO4
2-(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
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Lead Acid Accumulator
Charging: Anode(-) when discharging becomes
the cathode(-) when charging:PbSO4(s) + 2e Pb(s) + SO4
2-(aq)
Cathode(+) when discharging becomes the anode(+) when charging:
PbSO4(s)+ 2H2O(l) PbO2(s)+ 4H+(aq)+ SO4
2-
(aq)+2e 27
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
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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.
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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. 31
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)
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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
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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)
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