Electrochemistry

 

Ron Robertson

 

Applications of Redox Reactions

  • Terminology

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  • Zn (s) + Cu+2(aq) ® Zn+2 (aq) + Cu (s)

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  • This reaction is redox because electrons have been transferred.

  • Oxidation - loss of electrons, ox # increases, 
    Zn ® Zn+2

  • Reduction - gain of electrons, ox # decreases, 
    Cu+2 ® Cu

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  • Reducing agent - agent that allows reduction to occur, it is the substance that is oxidized and can provide electrons

  • Oxidizing agent - agent that allows oxidation to occur, it is the substance that is reduced and takes the electrons provided.

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  • Examples

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  • Corrosion

  • Combustion

  • Metabolism

  • Voltaic cells (batteries)

  • Electrolysis

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    Voltaic cells (Batteries)

  • Theory of operation

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  • Zn(s) + Cu+2(aq) ® Zn+2(aq) + Cu(s)

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  • This does not look like a battery because the electrons are being transferred directly between Cu+2 and the Zn strip. If we separate the reactions and force the electrons to travel through a circuit we could use these electrons as they try to get from the Zn to the Cu+2.

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  • The push of the electrons to go from Zn to Cu+2 is called the voltage.

  • Voltage = work/charge V = W/Q

  • Volt [=] Joule/coulomb

  • Anode - where oxidation takes place (- polarity)

  • Cathode - where reduction takes place (+ polarity)

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  • This push can be measured with a voltmeter.

  • Why do we need the salt bridge?

  • In the figure above the electrons from Zn ® Zn+2 will move through the wire and over to the other side to the Cu+2. Almost immediately the flow stops because of a charge buildup on both sides.

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  • Zn+2 ions are accumulating in the left compartment (+) and Cu+2 ions are leaving on the right side (-). We must allow ions to flow to equalize the charge. This is the function of the salt bridge.

  • With the salt bridge the circle is complete. Electrons flow through the wire from Zn to Cu+2 and anions move from the right side to the left and cations move from left side to the right to equalize the charge. This total movement of charge, ions in the internal circuit and electrons in the external, completes the voltaic cell.

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    Common Batteries

  • 1. Dry Cell (1.5 V)

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    Anode Zn ® Zn+2(aq) +2e

    Cathode 2NH4+ + 2e ® 2NH3(g) + H2 (g)

    The hydrogen gas is a problem so MnO2 is added; gaseous ammonia reacts with the Zn+2 ion.

  • 2. Alkaline (1.54 V)

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  • Anode Zn(s) + 2OH-(aq) ® ZnO(aq) + H2O + 2e

  • Cathode 2 MnO2(s) + H2O + 2e ® Mn2O3 (s) + 2OH-(aq)

  • no gases are formed

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  • no decline in voltage under high load
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  • 3. Lead storage (2.04V)

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  • Anode Pb° (s) + SO4-2(aq) ® PbSO4 (s) + 2e

  • Cathode PbO2(s) + 4 H3O+ + SO4-2(aq) ® PbSO4(s) + 6H2O

  • Net reaction

  • Pb + PbO2 + 2H2SO4 ® 2PbSO4 + 2H2O

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    4. Nickel/Cadmium (1.3 V)

  • Anode Cd(s) + 2OH-(aq) ® Cd(OH)2(s) + 2e

  • Cathode NiO2 (s) + 2H2O + 2e ® Ni(OH)2(s) + 2OH-(aq)

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  • can be recharged due to the fact that products are insoluble hydroxides that stay at the electrode surface
  • this battery produces a constant voltage until completely discharged
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  • Fuel Cell

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  • a special type of voltaic cell in which the reactants are continually supplied from an external source; the best known is the hydrogen / oxygen fuel cell used in the space program

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    Corrosion

  • Defined as the oxidation of a metal that results in a loss of structural strength and that is a result of exposure to the environment.

  • The most common is the corrosion of iron (rusting). Many metals oxidize but only those which form scaly oxidize coatings corrode. For example aluminum oxidizes readily but does not corrode easily.

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    Electrolysis

  • This is the opposite process from the galvanic cell. We use an electric current to force a nonspontaneous reaction.

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    Cell

    Electrode

    Function

    Polarity

    Battery

    Anode

    oxidation

    negative
     

    Cathode

    reduction

    positive

    Electrolysis

    Anode

    oxidation

    positive
     

    Cathode

    reduction

    negative

  •  Applications

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    1. Production of elements such as sodium and chlorine

    2. Plating of metals over another metal (or suitably prepared substance) for protection or beauty