May 20, 2014
Unit 8: Redox and Electrochemistry
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May 20, 2014
Oxidation Number
• numbers assigned to atoms that allow us to keep
track of electrons.
Rule #1: Oxidation number of any uncombined
atom is zero.
Example: C, H2, Al, Cl2...etc.
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Rule #2: The oxidation number of a monatomic
ion is equal to its charge.
Example: Na+, Mg2+, Cl-, S2-
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Rule #3: The oxidation number of the more
electronegative atom in a molecule or complex
ion is the same as the charge it would have if it
were an ion.
Example: NH3
of -3.
Nitrogen has oxidation number
Rule #4: The oxidation number of fluorine in a
compound always -1.
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Rule #5: Oxygen has an oxidation number of -2
in most compounds.
Exception 1: In peroxides (like H2O2), oxidation
number = -1.
Exception 2: When bonded to fluorine, oxidation
number = +2
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Rule #6: The oxidation number of hydrogen in
most of its compounds is +1 except when
bonded to metals, where it is -1.
Example: H2O, MgH2
Rule #7: In compounds, the elements of groups
1 and 2, and aluminum have oxidation numbers
+1, +2, and +3 respectively.
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Rule #8: The sum of the oxidation numbers in a
neutral compound is zero.
Example: NaCl, CaBr2, CCl4
Rule #9: The sum of the oxidation numbers in a
polyatomic ion is equal to the charge of the ion.
Example: SO32-, OH-
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*side note:
Charges are written with signs after the number:
2-, 3-, 2+, 3+
Oxidation numbers are written with signs before
the number
-3, -2, -1, +1, +2, +3,
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Example 1: Assign oxidation numbers
a) F2
b) Na2O
c) Fd) BH3
e) NaOH
f) PO43-
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Oxidation-Reduction Reactions
• also called Redox reactions
• reaction in which one or more electrons are
transferred from one atom to another
• Oxidation and reduction always happens together
Oxidation = losing electrons (oxidation
number increases)
Reduction = gaining electrons (oxidation
number decreases)
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reduction
electrons
gaining
losing
electrons
oxidation
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Why does redox happen?
• Atoms transfer electrons to another atom.
• The more electronegative atom attracts electrons
more strongly, resulting in a transfer of electrons.
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Example 2: Identify the following as oxidation or
reduction
a) I2 + 2e-
b) K
c) Fe2+
2I-
K+ + e-
Fe3+ + e-
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Oxidation numbers in Redox Reaction
1. Assign oxidation numbers to all elements
2. When an atom is oxidized, its oxidation #
increases
3. When an atom is reduced, its oxidation #
decreases
2K + Br2
2KBr
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Example 3:
Cu + AgNO3
Ag + CuNO3
Assign oxidation numbers.
• Which atom is oxidized?
• Which atom is reduced?
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Example 4:
2KBr + Cl2
2KCl + Br2
Assign oxidation numbers.
• Which atom is oxidized?
• Which atom is reduced?
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Example 5:
CH4 + 2O2
CO2 + 2H2O
Assign oxidation numbers.
• Which atom is oxidized?
• Which atom is reduced?
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Balancing Redox Reactions
• oxidation = reduction
# of electrons lost = # electrons gained
We will learn to balance redox reactions using
the half-reaction method.
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Half-reactions
• Equations that have electrons as reactants or
products
• One half reaction represents oxidation
• One half reaction represents reduction
Example 6:
SnCl4 + Fe
SnCl2 + FeCl3
Example 7:
Fe + CuSO4
Cu + Fe2(SO4)3
May 20, 2014
Using Half-Reactions to Balance Redox Equations
1. Identify the species oxidized and the species
reduced
2. Write the half-reaction
3. Multiply the half-reaction by the smallest
coefficient possible so that the # of e- is the
same.
4. Rewrite as a complete balanced equation (add 2
half-reactions together).
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Example 8:
Rewrite half reactions from example 6.
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Example 9:
Rewrite half reactions from example 7.
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Example 10: Balance the following using the halfreaction method.
H2S + Cl2
HCl + S
a. write half reactions
b. balance oxidation/reduction
c. rewrite balanced equation
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Electrochemistry
• Study of how chemical energy is converted to
electrical energy or vice versa.
Electrochemistry
Redox
Electrochemical
Cell
Voltaic Cell
Electrolytic Cell
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Voltaic Cells
• converts chemical energy to electrical energy
• spontaneous redox reaction
• generates a current
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Voltaic Cells
• consists of two half-cells
• Separate oxidation and reduction reaction
• Each half-cell contains:
> electrode
> solution
• Anode: oxidation
• Cathode: reduction
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Voltaic Cells
• Half-cells are connected by a salt bridge
> allows ions to pass from one side to another
> prevents build up of ions that prevent redox
reactions
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Example 11: Sketch the diagram from the animation
and identify
• anode and half-reaction
• cathode and half-reaction
• write the overall balanced cell reaction
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Cell Notation
• shows you the oxidation and reduction half-cells in
a voltaic cell
Anode
Zn
Zn2+
Oxidation
Cathode
Cu2+ Cu
Reduction
Salt Bridge
Example 12: Write the cell notation for the voltaic
cell in example 11.
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Electrochemical Cell Potential
• reduction potential: tendency of a substance to
gain electrons
> reduction potential of an electrode is measured
in volts
> standard reduction potential (E0)
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• Electrochemical cell potential: difference in
potential between the half-reactions
> potential must be > 0 in order for the redox
reaction to be spontaneous
> In a voltaic cell, the half-reaction with the lower
reduction potential will be the oxidation reaction
(opposite reaction given on chart)
E0cell = E0reduction - E0oxidation
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Example 12: Write the cell notation and calculate the
cell potential for the following redox reaction.
Sn(s) + 2Cu+(aq)
Sn2+ + 2Cu(s)
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Example 13: Write the cell notation and calculate the
cell potential for the following redox reaction.
Mg(s) + Pb2+(aq)
Pb(s) + Mg2+(aq)
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Example 14: Given the pair of half-reactions,
• write the balanced equation for the overall cell
reaction
• calculate the cell potential
• write the cell notation
Co2+(aq) + 2e-
Co(s)
Cr3+(aq) + 3e-
Cr(s)
May 20, 2014
Example 15: Given the pair of half-reactions,
• write the balanced equation for the overall cell
reaction
• calculate the cell potential
• write the cell notation
Fe2+(aq) + 2eI2(s) + 2e-
Fe(s)
2I-(s)
May 20, 2014