Oxidation States

1
Oxidation States

• Oxidation states are used as a means to track the
  movement of electrons in a reaction. An atoms
  oxidation state is the charge it would have if
  all bonds in the molecule were 100 ionic.
• To determine oxidation states for atoms in a
  molecule or ion
• Draw the Lewis structure.
• Assume that any bond between atoms of the same
  element is 100 covalent so each atom gets one
  electron.
• Assume that any bond between atoms of different
  elements is 100 ionic i.e. all electrons in the
  bond are given to the more electronegative atom.
• The sum of the oxidation states for every atom in
  the molecule/ion must equal the net charge of the
  molecule/ion.

2
Oxidation States

• Determine the oxidation state of each atom in the
  sulfate ion.
• CH3CO2H - C is more electronegative than H.)
• Peroxydisulfate, S2O92-

3
Redox Reactions

• When an atom Loses Electrons in a reaction, its
  oxidation state _____________ and the compound is
  said to be Oxidized.
• When an atom Gains Electrons in a reaction, its
  oxidation state is _____________ and the compound
  is said to be Reduced.
• These processes must be concurrent.

Some of you may have learned an Alberta version
of this mnemonic OIL RIG (Oxidation Is Loss
Reduction is Gain) Either works.
4
Redox Reactions

• Consider a simple redox reaction
• The oxidation state of the hydrogen atoms goes
  from ___ to ___.
• The oxidation state of the oxygen atoms goes from
  ___ to ___.
• Thus, the hydrogen atoms are ___________ and the
  oxygen atoms are ______________.

2 H2(g) O2(g) ? 2 H2O(g)

• When copper roofs tarnish, they turn green (like
  the Parliament buildings or Statue of Liberty).
  This is because the copper reacts with oxygen,
  carbon dioxide and water in the atmosphere to
  produce a green solid, Cu2(OH)2CO3
• _____ is oxidized. ____ is reduced.
• _____ is the reducing agent. ____ is the
  oxidizing agent.

2 Cu(s) O2(g) CO2(g) H2O(g) ?
Cu2(OH)2CO3(s)
5
Half Reactions

• Most redox reactions can be separated into two
  half-reactions
• In the oxidation half-reaction, electrons are
  products.
• In the reduction half-reaction, electrons are
  reactants.
• In the CHEM 1000 lab, you used zinc metal to
  precipitate copper metal from a clear blue
  solution of copper(II) sulfate. The net ionic
  equation for this reaction was
• Half reactions

6
Balancing Redox Reactions

• Balancing redox reactions is generally easiest if
  you divide the reaction into the two
  half-reactions, balance each of those then
  recombine them, making sure that all of the
  electrons cancel.
• In acid
• 1) Separate the half-reactions and balance all
  elements except O H.
• 2) Balance oxygen in each half-reaction by adding
  H2O(l).
• 3) Balance hydrogen in each half-reaction by
  adding H(aq).
• 4) Balance the charge in each half-reaction by
  adding electrons.
• 5) Multiply each half-reaction by whatever
  coefficient will make the electrons cancel when
  the two half-reactions are added.
• 6) Add the two half-reactions, canceling
  electrons (and possibly H2O(l) and/or H(aq)) as
  necessary.
• In base, follow the above steps then add enough
  OH-(aq) to exactly cancel all H(aq). By
  definition, a basic solution never has a
  significant concentration of H(aq).
• In neutral conditions, H(aq) is acceptable as a
  product however, OH-(aq) should be added to
  neutralize any H(aq) on the reactant side of the
  equation.

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• Balance these half-reactions. Tell whether the
  reactant is an oxidizing or reducing agent and if
  the overall process is an oxidation or reduction.
  Unless noted otherwise, all are carried out in
  acid solution, meaning that H or H and H2O may
  be used to balance the equation.
• (a) Br2(l) ? Br(aq)
• (b) VO2(aq) ? V 3(aq)

8

• 2. These half-reactions are in basic solution.
  You may need to use OH or the OH / H2O pair to
  balance the equation.
• (a) CrO2 (aq) ? CrO42 (aq)
• (b) Ni(OH)2(s) ? NiO2(s)

9
Balancing Redox Reactions

• The reactions here are in acid solution, meaning
  that H or H and H2O may be used to balance the
  equation.
• (a) MnO4 (aq) HSO3 (aq) ?
  Mn2 (aq) SO42 (aq)
• (b) Cr2O72 (aq) Fe2 (aq) ?
  Cr3 (aq) Fe3 (aq)

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• 4. The reactions here are in basic solution. You
  may need to add OH or the OH / H2O pair to
  balance the equation.
• (a) Fe(OH)2(s) CrO42 (aq) ?
  Fe2O3(s) Cr(OH)4 (aq)
• (b) PbO2(s) Cl (aq) ? ClO
  (aq) Pb(OH)3 (aq)