Chapter 18: Redox Reactions and Electrochemistry

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Chapter 18 Redox Reactions and Electrochemistry

• Oxidation-Reduction
• Oxidation States
• The Half Reaction Method for Balancing Redox
  Reactions
• Introduction to Electrochemistry
• Galvanic Cells
• Electrolytic Cells

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Example 1

• For each of the following reactions, identify
• the oxidizing agent and the reducing agent.
• 2 Al(s) 3 Br2(l) ? 2 AlBr3(s)
• 4 Fe(s) 3 O2(g) ? 2 Fe2O3(s)

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Rules for Assigning Oxidation Numbers

• Atoms in elements are assigned an oxidation
  number of 0.
• Monatomic ions are assigned an oxidation number
  equal to their charge.
• Oxygen has an oxidation number of 2 in all its
  compounds except when it occurs as peroxide ion
  (O22-), where it is 1.
• Hydrogen has an oxidation number of 1 when it is
  in a compound with nonmetals, and an oxidation
  number of 1 when combined with metals.
• In binary compounds, the more electronegative
  element is assigned a negative oxidation number
  equal to the charge it has when it is an ion.
• The sum of all the oxidation numbers of the atoms
  in a molecule is 0.
• The sum of all the oxidation numbers of the atoms
  in a polyatomic ion is the charge on the ion.

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Example 2

• Assign oxidation states to each atom in the
• following compounds
• CaCl2
• K3PO4
• C3H8
• K2Cr2O7

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Example 3

• Identify the oxidizing agent and the reducing
• agent in the following chemical reaction
• CH4(s) 3 O2(g) ? CO2(g) 2 H2O(g)

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• The Half Reaction Method for Balancing Redox
  Reactions
• Write two half-reactions, one for any oxidations
  occurring in the overall reaction, and one for
  any reductions.
• Balance each half-reaction as follows
• Balance all elements other than oxygen and
  hydrogen.
• Balance oxygen by adding the appropriate number
  of water molecules (H2O) to the side of the
  equation that needs more oxygen atoms.
• Balance hydrogen by adding the appropriate number
  of hydrogen ions (H) to the side of the equation
  that needs more hydrogen atoms.
• Balance the charge by adding the appropriate
  number of electrons to the side of the equation
  with the greater overall positive charge.
• Multiply each half-reaction by a whole number so
  that the electrons lost in the oxidation
  half-reaction equal the number of electrons
  gained in the reduction half-reaction.
• Add the two half reactions together, keeping all
  of the reactants together on the left side of the
  reaction arrow and all of the products together
  on the right side of the reaction arrow. The
  electrons will cancel so they are not shown in
  the final equation.
• Cancel any substances that appear on both sides
  of the equation. Check to make sure that the
  equation is balanced.

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Example 10

• Balance the following redox reactions that occur
• in acidic solution
• H2S(aq) HNO3(aq) ? S8(s) NO(g)
• Cr2O7 2-(aq) I-(aq) ? Cr3(aq) IO3-(aq)
• IO3- (aq) I- (aq) ? I2(aq)

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Figure 17.6 Cartoon of atoms reacting
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Figure 17.14 Common dry cell battery
KOH