Chapter 17: Oxidation and Reduction

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• Chapter 17 Oxidation and Reduction

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Foods and fuels are high in energy.

• Foods like glucose, C6H12O6, a simple sugar are a
  reduced form of matter. Glucose is oxidized when
  it reacts with oxygen by combustion or
  metabolism
• C6H12O6 6O2 6CO2 6H2O energy
• Note that this reaction is exothermic.
• CO2 and H2O are oxidized forms of matter.
• They are low in energy.

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Photosynthesis

• The reverse of the reaction on the previous slide
  is a reduction
• 6CO2 6H2O energy C6H12O6 6O2
• It is an endothermic reaction.
• Energy is added by sunlight.

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Oxidation Numbers(review)

• Oxidation Number-Ion charges and apparent
  charges assigned to atoms within compounds.
• Also called Oxidation State.
• Bookkeeping system for keeping track of
  electrons.
• Transition Elements may have several different
  oxidation numbers.

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Rules for Assigning Oxidation Numbers

• Uncombined elements have an oxidation number of
  zero K, Fe, H2, O2.
• For a compound, the sum of oxidation numbers is
  zero.
• For a polyatomic ion, the sum of oxidation
  numbers is equal to the charge on the ion.
• For a monatomic ion, the charge is the oxidation
  number Na is 1, O2- is -2.

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Oxidation Number Rules (continued)

• When oxygen is present in a compound or
  polyatomic ion, it is assigned an oxidation
  number of -2 (except for peroxides like H2O2,
  where it is -1).
• Hydrogen usually has an oxidation number of 1,
  except in metal hydrides like LiAlH4.

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Steps in Determining Oxidation Numbers

• Write down the known oxidation numbers in the
  formula. Set the missing Oxidation number to x.
• Multiply the oxidation number by the subscript.
• Write an equation where the sum of the oxidation
  numbers equal zero for a compound or the charge
  for a polyatomic ion.
• Solve for the missing oxidation number.

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Example Problems

• Determine the oxidation number of the elements
  in H2SO4.
• The oxidation number of H is 1 since it is
  combined with non-metals.
• The oxidation number of O is -2 since this is not
  a peroxide.
• The sum of the oxidation numbers is 0
• Solving for x
• 0 2(1) x 4(-2) 0 -6 x x 6

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Cr2O72-

• Determine the oxidation number of Cr in Cr2O72-.
  
• The oxidation number of O is -2. The sum of the
  oxidation numbers is -2.
• Set up an equation, let x be the oxidation number
  of Cr.
• -2 2x 7(-2) 2x - 14
• 2x -2 14, 2x 12 x 12/2 6

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REDOX

• Many reactions are oxidation-reduction reactions.
  These are sometimes termed REDOX reactions.
• Most of the time, oxidation and reduction occur
  in the same reaction. One of the reactants is
  oxidized, the other is reduced.
• REDOX reactions involve the transfer of electrons
  from one reactant to another.
• There are several definitions of oxidation and
  reduction.

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Oxidation

• The term oxidation may be defined several
  different ways
• 1. Gain of oxygen
• 2. Loss of hydrogen
• 3. Loss of electrons
• 4. Increase in oxidation number.

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Example Oxidation Reactions

• Combustion
• C O2 CO2
• CH4 2O2 CO2 H2O heat
• Corrosion (rusting)
• 4Fe 3O2 2Fe2O3
• Loss of hydrogen
• CH3OH CH2O H2
• Loss of electrons (increase in oxidation number)
• Mg Cl2 Mg2 2Cl-

Cu
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Reduction

• The term reduction is also defined in several
  ways
• 1. Loss of oxygen
• 2. Gain of hydrogen
• 3. Gain of electrons
• 4. Decrease in Oxidation Number

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Reduction Reactions

• Reduction of metal oxides
• CuO H2 Cu H2O
• Loss of oxygen
• 2KClO3 2KCl 3O2
• Gain of hydrogen
• CO 2H2 CH3OH
• Gain of electrons
• Cu2 2e- Cu

heat
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A REDOX Reaction

• CuO H2 Cu H2O
• In this reaction CuO is reduced (the oxidation
  number of Cu goes from 2 to O).
• H2 causes the reduction and is called the
  reducing agent.
• H2 is oxidized (The oxidation number goes from 0
  to 2).
• CuO caused the oxidation and is called the
  oxidizing agent.

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Definitions

• Oxidizing Agent The reactant in a chemical
  reaction which causes another reactant to be
  oxidized, and is itself reduced.
• Reducing Agent The reactant in a chemical
  reaction which causes another reactant to be
  reduced, and is itself oxidized.

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Oxidation/Reduction Definitions
Insert figure 17.5
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Oxidizing Agents

• Oxygen, O2
• 4CH4 3O2 2C2H2 6H2O
• Permanganate ion, MnO4-
• MnO4- 2Fe2 8H Mn2 5Fe3 4H2O
• Dichromate ion, Cr2O72-
• 8HCr2O72-3C2H5OH 2Cr33C2H4O7H2O
• Chlorine, Cl2
• Mg Cl2 Mg2 2Cl-

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Example REDOX Reaction

• Consider the reaction
• Cu(s) 4H(aq) 2NO3-(aq) Cu2(aq)
  2H2O(l) NO2(g)
• Assign oxidation numbers to each element in the
  reactants and products. Which elements are
  oxidized? Which elements are reduced? What is
  the oxidizing agent? What is the reducing agent?

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Answer

• Cu(s) 4H(aq) 2NO3-(aq) Cu2(aq)
  2H2O(l) NO2(g)
• Oxidation Numbers
• Reactants Products Process Agent
• Cu 0 Cu2 2 ox reducing
  agent
• H 1 H 1 nc
• O -2 O -2 nc
• N -1 N 3(-2) N 0 N 2(-2)
• N 5 N 4 red
• NO3- oxidizing agent

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Reducing Agents

• Carbon, C
• SnO2 C Sn CO2
• Hydrogen, H2
• WO3 3H2 W 3H2O
• Carbon monoxide, CO
• FeO CO Fe CO2

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Half Reactions

• REDOX reactions can be divided into two parts
  reduction half-reaction and oxidation
  half-reaction.
• Dividing equations into half reactions is useful
  in balancing REDOX reaction equation.

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Balance the following equation

• MnO4- Fe2 Fe3 Mn2
• Split the equation into two half reactions
• MnO4- Mn2
• Fe2 Fe3
• Balance atoms using coefficients (done)
• Balance O using H2O
• MnO4- Mn2 4H2O
• (continued)

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Balance the following equation(continued)

• MnO4- Fe2 Fe3 Mn2
• Balance the hydrogens with H
• 8H MnO4- Mn2 4H2O
• Balance electrons with e-
• Fe2 Fe3 e-
• 5e- 8H MnO4- Mn2 4H2O
• Multiply by whole numbers to get equal electrons
• 5Fe2 5Fe3 5e-
• (continued)

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Balance the following equation(continued)

• MnO4- Fe2 Fe3 Mn2
• Add equations together
• 5e- 8H MnO4- Mn2 4H2O
• 5Fe2 5Fe3 5e-

5e- MnO4- 5Fe2 8H Mn2
5Fe3 4H2O 5e-
Cancel Electrons and others which are on both
sides.
MnO4- 5Fe2 8H Mn2 5Fe3
4H2O
Check to make sure all atoms and charges are
balanced.
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Electrolytic Cells

• Electricity (DC current) can be used to cause
  REDOX reactions to occur. This process is called
  electrolysis.
• Molten NaCl can be separated into it elements
• 2NaCl(l) 2Na(l) Cl2(g)
• reduction Na e- Na (cathode)
• oxidation 2Cl- Cl2 2e- (anode)

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Electrolytic Cells
Insert figures 17.7 and 17.8
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Electroplating

• A metal can be plated onto a cathode from a
  solution of the ions of the metal.
• The ions are replaced by oxidation at the anode
• cathode reaction Ag(aq) e- Ag(s)
• anode reaction Ag(s) Ag(aq)

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Electroplating
Insert figure 17.9
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Electrochemical Cells

• A chemical system that uses an electric current
  to produce a chemical reaction or that generates
  electricity as a result of a chemical reaction is
  called an electrochemical cell.
• Electrochemical cells that use electricity to
  make a chemical reaction occur are called
  electrolytic cells.

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Voltaic Cells

• Chemical reactions can produce electricity.
• A chemical system that generates electric current
  is called a voltaic cell or a galvanic cell.
• An example of a chemical reaction that can
  generate electricity is the reaction of zinc
  metal with copper (II) ions
• Zn(s) Cu2(aq) Zn2(aq) Cu(s)

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Zn(s) Cu2(aq) Zn2(aq) Cu(s)

• When zinc metal is placed in a Cu2 solution, the
  zinc dissolves to form Zn2 ions, an the copper
  metal plates out on the surface of the zinc.
• The reaction is a REDOX reaction
• Oxidation Zn(s) Zn2(aq) 2e-
• Reduction Cu2 2e- Cu(s)
• These are called half reactions.

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Zn(s) Cu2(aq) Zn2(aq) Cu(s)

• The two half reactions
• Oxidation Zn(s) Zn2(aq) 2e-
• Reduction Cu2 2e- Cu(s)
• The generate an electrical voltage two half cells
  are constructed
• Zinc metal in a solution of Zn2 ions
• Copper metal in a solution of Cu2 ions
• The metals are connected by a wire through which
  the electric current can flow.
• The two solutions are separated by a porous
  barrier or connected by a salt bridge.

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Insert figure 17.10, 17.11
Voltaic Cells
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Half cell reactions

• In the solution containing the zinc, oxidation
  occurs
• Zn(s) Zn2(aq) 2e-
• The zinc is called the anode.
• In the solution containing the copper, reduction
  occurs
• Reduction Cu2 2e- Cu(s)
• The copper is called the cathode.
• Electrons flow through the wire from the anode to
  the cathode.
• Charges in the solutions are balanced by ion
  migration through the salt bridge or porous
  barrier.

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Voltaic cells

• Commercial batteries operate on a principle
  similar the the Cu/Zn cell.
• There are many electrochemical cells available
  today.
• Two examples are the electric flashlight battery
  or dry cell and the lead/acid storage battery
  used in motor vehicles.
• The corrosion of metals is another example of en
  electrochemical cell. Metal ions are oxidized
  and oxygen is reduced.