Transcript Unit 16 REDOX

(Redox)
 1.
Synthesis
 2. Decomposition
 3. Single Replacement
 4. Double Replacement
 * Combustion
 assigned
to atoms or ions
as a way to keep track of
electron transfers
The oxidation number for each atom
in free elements is always zero (ex:
H2, Na, S8)
 The oxidation numbers of ions are the
same as the charge on the ion.
In MgCl2 the Mg+2 ion has an
oxidation number of +2. Each of the
two Cl- ions has an oxidation number
of –1.

 Group
1 metals always have a
+1 oxidation number.
 Group 2 metals always have a
+2 oxidation number.
 Oxygen
has a –2 oxidation
number in practically all of its
compounds. The exceptions are
in peroxides, such as H2O2 and
Na2O2 where the oxidation
number is –1 and in compounds
with fluorine (OF2) in which it is
+2.
 Hydrogen
has an oxidation
number of +1 in all its
compounds except metal
hydrides formed with Group 1
and Group 2 metals, which it
is –1.
 The
sum of the oxidation
numbers in a compound must be
zero.
 The sum of the oxidation
numbers in a polyatomic ion must
be equal to the charge on the ion.
In the CO32- ion, the three
oxygens provide a total of –6; the
carbon must then be +4.
 Situation
1:
Find the oxidation numbers of
N and O in N2O.
 Situation
2:
Find the oxidation numbers of
K, Mn and O in KMnO4.
(next slide for chart)
Element K
Each
Total
Mn
O4
 reactions
in which both
oxidation and reduction
occur
 the
loss of electrons by an
atom or an ion, increasing the
oxidation number
 Ex: Na0
Na+1(Na loss 1e-)
 Na
Na+1 + 1e Ex:
O-2
O0 (O loss 2e-)
 the
gain of electrons by an
atom or an ion, reducing the
oxidation number
 Ex: F0
F-1 (F0 gain 1e-)
 F0 + 1eF-1
 Ex:
Ca+2
Ca0 (Ca+2 gained 2e-)
 LEO
(Lose Electrons Oxidations)
 oxidation # increases

oxidation # decreases
 GER
(Gain Electrons Reduction)
 LEO
the lion says GER!
◦ Electronic equations or equations
in which only the atoms/ions
being oxidized or reduced are
shown.

 Ex:
2Al0 + 3Cl20 = 2Al+3Cl3-1
 Oxidation
½:
Al0 → Al+3 + 3e-
 Reduction
½: Cl20 + 2e- → 2Cl-1
 Must
conserve the number of
electrons!
 Number of electrons lost
equals number of electrons
gained
Oxidizing Agent: causes something
else to undergo oxidation by taking
on its electrons. This is the
substance reduced.
Reducing Agent: causes something
else to undergo reduction by giving
its electrons. This is the substance
oxidized.
 ions
that are not changed
during a redox reaction
Identifying Redox Reactions
 Once oxidation numbers are assigned,
the atom that has shown an increase
can be identified as the one that has
undergone oxidation. The atom that
has a decrease in oxidation number
can be identified as the one that has
undergone reduction.

 *Double
Replacement
reactions are never redox
reactions because NOTHING
CHANGES!
 Ex:
HCl + NaOH = NaCl + H2O

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Table J – Reactivity Series of Selected Metals
and Nonmetals
Lithium will react more readily with a
nonmetal than will any other metal on the
list.
Hydrogen (not a metal!!) is shown on the
metal side to illustrate what metals will react
with acids and which will not. All metals
above hydrogen will react spontaneously with
acids, while those below hydrogen will not.
 if
a reaction proceeds to
completion without adding energy
sources once it is started, then it
is spontaneous
You must have something to oxidize
and something to reduce.
 Mg(s) and Al(s) can only be oxidized
 Mg+2 (aq) and Al+3(aq) can only be
reduced

*One metal must be neutral, the
other an ion!!!
 Cu(s)
+ Al(s) = no rxn both
cannot be oxidized
 2Cr+3 + 3Mn(s) = 3Mn+2 + 2Cr(s)
 Cr+3 + Ni(s) = no rxn
 Ni is lower than Cr on Table J,
must be higher.
Again, there must be something
to oxidize and something to
reduce:
 F2 and Cl2 can only be reduced
 F- and Cl- can only be oxidized
 Cl2
+ Br2 = no rxn – can’t both be
reduced
 Cl2 + 2Br- = 2Cl- + Br2
 I2 + Br- = no rxn
 I is lower on the table than Br,
must be higher
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HCl + Au(s) = Au will not react with H
HCl + Sr(s) = SrCl2 + H2
Yes because Sr is more likely to oxidize than
H2
2Na(s) + ZnCl2(aq) = NaCl + Zn
Yes because Na is higher on Table J so it will
oxidize while Zn reduces.
Voltaic Cells: ½
cells of 2 different metals connected by wires –
produces electricity.

redox is spontaneous.
Electrons flow from the metal being oxidized
(reducing agent) to the ion being reduced
(oxidizing agent).

Salt Bridge:
 These are batteries which
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functions to complete the circuit by allowing ions to flow
from reducing side to oxidizing side.
Diagram:
Which reaction is occurring in the cell?
Zn(s) + Cu+2 = Zn+2 + Cu(s)
or
Zn+2 + Cu(s) = Zn(s) + Cu+2
The anode is Zn0 and the cathode is Cu+2
lectrochemical Cells
electricity is used to produce a nonspontaneous redox
reaction.
 Electricity is required! This is not spontaneous!! Requires the use of
a Battery.
 The anode is connected to the  and the cathode is connected to the 
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Examples: Electroplating
Diagram:
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*The only way to get a group one metal in its elemental
form is by electrolysis of their fused salts.
Example: KBr = K(s) + Br2(s)