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```							Chapter 20
Oxidation Numbers

1.   Elements always have oxidation number = 0.
2.   Column I alkalai metals in compounds always
have oxidation number = +1.
3.   Column II alkaline earth metals in compounds
always have oxidation number = +2.
4.   Aluminum and gallium oxidation numbers are +3,
zinc and cadmium are +2, and silver is +1.
5.   Hydrogen normally has oxidation number = +1 in
compounds except when combined with Colunm I
or Column II elements; then rules 2 and 3 apply.
6.   Oxygen normally has oxidation number = -2
except in H2O2 (rule 4 has priority), or Column I
and Column II oxides, where rules 2 and 3 apply.
7.   To calculate other oxidation numbers pretend
oxidation numbers are per atom charges and
make all “charges” on all atoms total up to overall
charge on ion or molecule containing atoms.
Calculating Oxidation Numbers

   Each oxide ion has a charge of -2
   7 oxide ions have a subtotal charge of -2 x 7 = -14
   Since the formula has to be uncharged the 2 manganese ions have to have a +14 subtotal
   The +14 subtotal divided evenly over 2 manganese ions gives each manganese +14 / 2 = +7
   This compound is manganese(VII) oxide
   Work oxidation numbers of Cr and S in Cr2(SO4)3 (Hint: treat SO42- as a single particle)
Oxidation and Reduction

1.   In simple chemical reduction-oxidation (redox)
reactions one reactant substance contains an
atom whose oxidation number increases when
product is created. This substance becomes
oxidized in the reaction. This substance is called
a reducing agent or reductant because it causes
another substance to become reduced.

2.   The other reactant substance contains an atom
whose oxidation number decreases when product
is created. This substance becomes reduced in
the reaction.    This substance is called an
oxidizing agent or oxidant because it causes
another substance to become oxidized.
Oxidation -Reduction Reactions I

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Oxidation - Reduction Reactions II

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Common Oxidation States

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Balancing Redox Reactions
An alternative to textbook method called the REDOHX
method is outlined below.

1.   Balance Reducing and oxidizing atoms first.
2.   Update total “charges” on reducing and oxidizing
atoms by multiplying oxidation numbers by all
appropriate subscripts and coefficients.
3.   Balance Electrons by adjusting coefficients in
front of oxidizing and reducing agents.
4.   Balance atoms which Don’t fit other categories.
5.   Balance Oxygens by adding H2O.
6.   Balance H by adding H+
7.   Xtra work only when balancing in base solution.
-
Add OH to both sides to destroy H+ and then
eliminate redundant H2O.
Cu 2+         +     Fe 0      Redox Reaction

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CuO + C Redox Reaction

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Mercury (II) Oxide Decomposition

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SnCl2 + Zn0 Redox Reaction

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Standard Cell Potentials
A standard reduction potential (previous slide) is a
stoichiometry-independent potential energy difference
measured in voltage units associated with a reduction
reaction done under standard conditions.

An oxidation reaction is simply the reverse of a
reduction reaction and therefore has the same voltage
with the sign reversed as the voltage of the (opposite)
reduction reaction found in a reduction potential table.

The total cell potential for a redox reaction done under
standard conditions is sum of the voltages (energies)
for the oxidation part plus the reduction part.

Remember to change the sign of voltage for oxidation
reaction before adding voltage for reduction reaction.
Relationship Between E° and G°
To convert an an energy in voltage units into a Gibbs
free energy three things need to be done:

1.   Voltages are stoichiometry-independent (G° is
not). Need to balance redox reaction to figure out
how many electrons involved in redox reaction
(“e”) and multiply by this number to fix this.
2.   Sign convention of voltages are opposite that of
G°. Need to change sign of energy.
3.   Voltages based on coulomb quantities and G°
based on mole quantities. Need to multiply by
Faraday’s constant (F = 96,500 C/mol) to fix this.

In Summary: G° = -eFE°
Nonstandard E Values
Nernst Equation: Analogous to nonstandard ∆G equation
∆G = ∆Go + RTlnQ
E = Eo - (RT/eF)lnQ
• Notice effect of opposite sign convention on direction of
deviation from standard value
• Notice RT (kJ/mol) becomes RT/eF (J/coul)
• R = 0.008314 kJ/mol-K (∆G) vs. R = 8.314 J/mol-K (E)
Nonstandard E Values
Batteries in real world seldom have standard conc’s of
all redox reactants and products. To figure out if
voltage higher or lower than calculated for standard
conditions use Le Châtelier’s principle to decide if
reaction more or less spontaneous than standard and
interpret this in terms of more or less positive E value.

Problem 20.58:

Al(s) + 3 Ag+(aq)  Al3+(aq) + 3 Ag(s)

(a) Dilute anode cell   (b) Increase Al(s)
(c) Increase AgNO3 vol. (d) Add HCl to AgNO3 cell
To convert between number of moles of electrolysis
product (n) and amps of current (I), or t (time in
seconds), or coulombs of charge (C = It), or number of
electrons involved in redox reaction (“e”) Faraday’s
law is used:
n = It/(eF)

Problem 20.86:

(a) What mass of Mg formed by passsing 5.25 A thru
molten MgCl2 for 2.50 days?
(b) How many minutes to make 10.00 g of Mg from
molten MgCl2 using I = 3.50 A?

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