Electricity from Chemical Reactions

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Electricity from Chemical Reactions
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Electrochemistry

• The production of electrical energy from chemical
  reactions
• Redox reactions involve the transfer of electrons
• Redox means that reduction and oxidation are
  occurring simultaneously

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Reduction

• Occurs when there is a decrease in oxidation
  number Zn2 ? Zn
• Gains electrons
• Loses Oxygen
• Converting a complex substance into a simpler
  form i.e. smelting iron to produce the pure metal
  iron

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Oxidation

• Occurs where there is an increase in oxidation
  number Zn ? Zn2
• Loses electrons
• Gains oxygen
• The reaction used to describe the reaction of any
  substance with oxygen

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Determining Oxidation Numbers

• The atoms in elements have an Oxidation Number
  of zero eg Fe, C, Cl2
• For a neutral molecule, the sum of the oxidation
  numbers are zero eg CO2
• For a monatomic ion, the oxidation number is the
  same as its charge Cl , Na

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Determining Oxidation Numbers

• Oxygen usually takes 2 in compounds. In
  peroxides (H2O2 BaO2) it is 1
• Hydrogen takes 1 in compounds, except in
  hydrides (NaH, CaH2) where it takes 1

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Determining Oxidation Numbers

• For a polyatomic ion, the sum of the oxidation
  numbers of its component atoms is the same as its
  charges
• For polyatomic molecules or ions, the, most
  electronegative element has a negative oxidation
  number and the least electronegative element has
  a positive oxidation number

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Redox Half Reactions

• Consider the reaction when a strip of zinc is
  dropped in a solution of Copper Sulphate
• Zn(s) Cu 2(aq) ? Zn2(aq) Cu(s)
• Electrons are transferred from zinc atoms to
  copper ions
• Reaction occurs spontaneously, that is with no
  external force or energy being applied

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Redox Half Reactions

• Redox reactions consist of two half reactions
• Oxidation Zn(s) ? Zn2(aq) 2e1
• Reduction Cu 2(aq) 2e1 ? Cu(s)
• It is possible to use redox reactions to produce
  electricity

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Galvanic Cells

• Also called Electrochemical Cells
• Achieved by separating the half equations into
  half cells
• Transferred electrons are forced to pass through
  an external circuit
• Such an apparatus is called a Galvanic Cell

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Galvanic Cells
Flow of electrons


zinc
copper
Salt bridge
Cu2
Zn2
Negative Electrode (ANODE)
Positive Electrode (CATHODE)
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Standard Electrode Potentials

• The electrical potential of a galvanic cell is
  the ability of the cell to produce an electric
  current.
• Electrical potential is measured in volts
• Cannot measure the electrode potential of an
  isolated half cell
• Can measure the difference in in potential
  between two connected half cells

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Standard Electrode Potentials

• Electrical potential of a cell results from
  competition between 2 half cells for electrons
• Half cell with the greatest tendency to attract
  electrons will undergo REDUCTION
• Other half cell will lose electrons and undergo
  OXIDATION

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Standard Electrode Potentials

• The Reduction Potential of a half cell is a
  measure of the tendency of the oxidant to accept
  electrons and so undergo reduction
• The difference between the reduction potentials
  of the two half cells is called the Cell
  Potential Difference

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Standard Electrode Potentials

• The Standard Cell Potential Difference (E0
  cell) is the measured cell potential difference
  when the concentration of each species 1M,
  pressure 1 atm and Temp 25 C
• E0 cell E0 oxidant E0 reductant

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Standard Electrode Potentials

• A Standard Hydrogen Half cell is used as a
  comparative measure the reduction potentials of
  other cells
• The SHE is given a value of 0.00 V
• All other half cells are given a reduction
  potential value in comparison to this SHE by
  being connected to it

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Standard Hydrogen Electrode
Platinum wire
Glass sleeve
H2 gas (1 Atm)
Salt Bridge to Other half-cell
1.00M Acid solution
Platinum electrode
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Standard Hydrogen Electrode

• SHE is used to measure reduction potential of
  other cells
• If a species accepts electrons more readily than
  hydrogen, its electrode potential is positive
• If a species accepts electrons less readily than
  hydrogen, its electrode potential is negative

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Electrochemical Series

• The reaction that is higher on the
  electrochemical series will occur as it appears
  and will reverse the direction of the reaction
  that occurs lower on the table

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Potential Difference

• Is measured by a volt meter
• Can be estimated by using electrochemical series
• Connect Mg2/Mg and Cl2/Cl half cells get a
  potential difference of 3.7V
• Looking at the electrochemical series

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Potential Difference

• Cl2 2e ? Cl has an E0 of 1.36V
• Mg2 2e ? Mg has an E0 of 2.38V
• The potential difference can be calculated
• 1.36 ( 2.38) 3.74V

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Galvanic Cells

• Primary Cells
• Produce energy until one component is used up,
  then discarded
• Secondary Cells
• Store energy and may be recharged

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Primary Cells

• Dry Cells
• Alkaline Cells
• Button Cells

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Dry Cells

• The ordinary zinc carbon cell
• Anode oxidation ()
• Zn (s) ? Zn 2 (aq) 2e
• Cathode oxidation ()
• 2MnO2 (s) NH4 (aq) 2e ? Mn2O2 (s) 2NH3
  (aq) H2O (l)

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Dry Cells

• The new cell produces about 1.5V
• Once reaction reaches equilibrium its flat

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Dry Cell
Metal Cap ()
Mixture of Carbon Manganese Dioxide
CathodeCarbon Rod
Ammonium Chloride Zinc Chloride Electrolyte
Anode Zinc Case ()
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Alkaline Cells

• The ordinary zinc carbon cell
• Anode oxidation ()
• Zn (s) ? Zn 2 (aq) 2e
• Immediately reacts with OH ions in the
  electrolyte to form zinc hydroxide
• Zn (s) 2OH (aq) ? Zn(OH)2 (s) 2e

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Alkaline Cells

• Cathode reduction ()
• 2MnO2 (s) H2O(l) 2e ? MnO2 (s) OH (aq)
  H2O (l)
• Five times the life of the dry cell

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Alkaline Cell
Metal Cap ()
Cathode outer steel case
Potassium Hydroxide Electrolyte
Powdered Zinc
AnodeSteel or Brass
Mixture of Carbon Manganese Dioxide
Metal Base ()
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Button Cells

• Used in very small applications like watches,
  cameras etc.
• Two main types
• Mercury zinc and silver zinc
• Anode Oxidation ()
• Zn (s) 2OH (aq) ? Zn(OH)2 (s) 2e

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Button Cells

• Cathode Reduction ()
• depends on the type of battery
• HgO(s) H2O (l) 2e ? Hg (l) 2OH (aq)
• Ag2O(s)H2O (l) 2e ? 2Ag (s) 2OH (aq)
• Produce an almost constant 1.35V

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Button Cell
Metal Cap ()
Zinc Powder
Cathode outer container of nickel or steel ()
Electrolyte
Mercury Oxide
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Secondary Cells

• Lead Acid (Car Battery)
• Nickel cadmium Cells
• Fuel Cells

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Lead Acid Battery

• Car Batteries p 211-2
• Also called storage batteries or accumulators
• Each cell produces 2 volts so typical 12 volt car
  battery contains 6 cells
• Both electrodes are lead plates separated by some
  porous material like cardboard

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Lead Acid Battery

• Positive electrode is coated with PbO2 Lead (IV)
  Oxide
• The electrolyte is a solution of 4M sulfuric acid

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Lead Acid Battery

• Anode Oxidation ()
• Pb(s) SO4 2- ? PbSO4 (s) 2e
• Cathode Reduction ()
• PbO2(s) SO4 2- 4H 2e ? PbSO4 (s) 2H2O
  (l)
• Overall Reaction
• Pb(s) PbO2(s) 2H2SO4 ? 2PbSO4 (s) 2H2O (l)

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Nickel Cadmium Cells

• Often called Nicads
• Electrodes are Nickel and Cadmium
• Electrolyte is Potassium Hydroxide
• Reactions involve the hydroxides of the two metals

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Nickel Cadmium Cells

• Anode (Oxidation) ( )
• Cd (s) 2OH (aq) ? Cd(OH)2 (s) 2 e
• Cathode (Reduction) ()
• NiO-OH (s) H2O (l) e ? Ni(OH)2 (s) OH
  (aq)
• Overall Reaction
• Cd (s) NiO-OH(s) H2O(l) ? Cd(OH)2 (s)
  Ni(OH)2 (s)

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Fuel Cells

• Limitation of dry cells looked at so far is that
  they contain reactants in small amounts and when
  they reach equilibrium.
• Primary Cells are then discarded, secondary cells
  are then recharged
• A cell that can be continually fed reactants
  would overcome this and allow for a continual
  supply of electricity

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Fuel Cells

• Fuel cells transform chemical energy directly
  into electrical energy
• 60 efficiency
• Space Program uses hydrogen and oxygen with an
  electrolyte of Potassium Hydroxide

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Fuel Cells

• Anode Oxidation ()
• H2(g) 2OH (aq) ? 2H2O (l) 2e
• Cathode Reduction ()
• O2(g) 2H2O(l) 4e ? 4OH(aq)
• Overall Equation
• H2(g) O2(g) ? 2H2O (l)

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Hydrogen Oxygen Fuel Cell


Electrolyte
Oxygen Gas Inlet
HydrogenGas Inlet
Porous Anode
Porous Cathode
Water outlet