Electrochemistry revision cards- Electrochemical cells, Electrolysis and Oxidation-Reduction Reactions

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  • Created by: Liv
  • Created on: 12-03-11 05:06


  • Relationship between electricity and chemical reactions

Two types of electrochemical processes: 

  • Generation of electric current from spontaneous chemical reactions
    • Electrochemical Cells
    • Exothermic
  • Use of electricity to cause non-spontaneous reactions to occur
    • Electrolysis
    • Endothermic
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Oxidation & Reduction


  • Loss of electrons
  • Gain of Oxygen
  • Loss of Hydrogen
  • Increase in Oxidation number


  •  Gain of electrons
  • Loss of Oxygen
  • Gain of Hydrogen
  • Decrease in Oxidation number
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Metal Displacement Reaction

  • Reactions between a metal & a solution containing ions of a different metal
  • Reaction will only occur if a reactive metal is place in a less reactive solution
    • use activity series to determine if reactions will occur


  • Metal dissolves and the ions of the other metal are reduced to elemental state and deposits out of the solution
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Rules for Oxidation States

  • Elemental State = 0
  • Monatomic Ions = charge on ion
  • The sum of oxidation states of atoms in a neautral molecule/ ionic compound= 0
  • The sum of oxidation states of atoms in a polyatomic ion = charge on ion
  • Group 1= +1
  • Group 2 = +2
  • Oxygen = -2
    • except peroxides = -1
    • and F20 = +2
  • Hydrogen = +1
    • except metal hydrides = -1
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Galvanic Cell

  • The Zinc anode becomes smaller as ions go into solution
  • Electrons flow from anode to cathode
  • Copper ions in solution become copper atoms which deposit on strip
  • Cations move from anode to cathode
  • Anions move from cathode to anode
  • Anode is oxidation
  • Cathode is reduction
  • Anode (-)
    • electrons are generated by oxidation
  • Cathode (+)
    • electrons are accpeted in reduction

Salt Bridge

  • filter paper containing an electrolytic solution ( KNO3 or NaNO3)
  • it completes the cicuit and allows ions to move between each half-cell
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Commercial Galvanic Cells

  • Galvanic cells are a source of direct current & provide portable electrcity

Classified into two groups:

  • Primary Cells
    • cannot be recharged
      • Lelanche Cell
      • Alkaline Dry Cell
  • Secondary Cells
    • can be recharged- supply electrical energy to reverse the cell reaction
      • Lead-acid Battery
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Primary Cells-Dry Cells

  • Dry cells have electrolytes in the form of pastes
  • Lelanche Cell


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Primary Cells-Dry Cells-Lelanche

Lelanche Cell-

  • The outer zinc casing is the anode
  • The graphite rod is surrounded by a paste containing manganese dioxide which is the cathode
  • Ammonium Chloride and Zinc Chloride are in the paste and act as the electrolyte (salt bridge)

Advantage: used in many applications


  • short shelf life (zinc and ammonium in paste react and cause leaking and deterioration in the cell)
  • Rapid drawing of current from the cell causes ammonia to buld up and this causes a drop in voltage
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Primary Cells-Dry Cells- Alkaline

  • Alkaline Dry Cell
    • Composition is similar to Lelanche cell-
    • except powered zinc anode is used 
    • electrolyte is 7mol/L of KOH


  • has a longer working life
  • supplies current more rapidly


  • is more expensive though


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Secondary Cell-Lead-acid

Lead-Acid Batteries are commonly used in motor vehicles

  • Electrodes in each cell consist of a bank of lead grids supporting a large surface area of the electrode material
  • Negative electrode (anode) grid is filled with spongy metallic lead
  • Positive electrode (cathode) grid is filled with brown lead oxide
  • Eletrolyte is sulfuric acid


  •  Reachargeable & Cheap


  • Enclosed in heavy casing-limits use in application
  • Lead can do body damage
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Secondary Cell-Lead-acid


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  • process where an electric current is used to drive a non spontaneous chemical reaction
  • electrical energy is converted to chemical energy
  • products have higher potential energy than reactants
  • current comes from power pack/battery
  • a cell which electrolysis occurs is called an electrolytic cell
  • electrons travel through external circuit


  • electroplating-improve appearance, reduce corrosin
  • electrorefining-higher levels of purity
  • extraction of reactive metals from their ores
  • manufacturing of aluminium, sodium hydroxide, chlorine, hydrogen
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Electrolysis- the electrodes & electrolyte


  • electrodes may/ may not take part in reaction
    • Inert electrodes do not take part in reaction
  • chemical reactions occur at interface between electrode and electrolyte
  • anode (+)
  • cathode (-)
  • battery pulls electrons from anode and pushes them to cathode


  • produces ions in the solution
  • ions may/may not take part in reaction
    • is water is present it is preferentially oxidised
    • or a reactive metal cathode (Zinc Chloride)
  • Cations to Anions
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Electrolysis of NaCl


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Electrode Potential


  • the nature of the element
  • the concentration of its ions in solution
  • the temperature of the solution

Amount of metal does not influence e.m.f. Temperature & concentration must be stated when comparing electrode potentials of different elements. Standard conditions are 298 K (25 degrees) and 1mol/dm3

  •  Increasing concentration increases the electrode potential
  • Temperature and concentration affect the voltage
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Predicting the Products

For Molten Salts

  • cations are reduced at cathode
  • anion is oxidised at anode

For Aqueous Solutions-at anode

  • Nitrate & Sulfate are never oxidised
  • Water is oxidised to produce Oxygen
  • Anions may be oxidised (depends on water)
  • Chloride is oxidised if concentrated
  • If a non-inert metal is used as electrode; it may oxidise & go into solution as ions
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Predicting the Products

For Aqueous Solutions- at cathode

  • Metal ions of less reactive metal (Cu& Ag) are reduced in preference to water
  • Water is reduced in preference to metal ions of reactive metals (Na, Mg, Al)

Anode Reaction:

  • More positive e.m.f value is oxidised (of the two competing)
  • Water is oxidised producing Oxygen gas

Cathode Reaction:

  • Ions with more positive e.m.f value is reduced (of the two competing)
  • Water is reduced producing Hydrogen gas
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Electrolysis of Water

  • Pure water does not conduct electricity-therefore KNO3 or Dilute Sulfuric acid is added.
  • Hydrogen gas is produced at the cathode
    • H ions are discharged-accepting electrons to form hydrogen molecules
  • Oxygen gas is produced at the anode
    • OH ions are discharged-form water and oxygen molecules


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