Transcript Ch20 Lesson20_3

20.3 Describing Redox Equations >
Chapter 20
Oxidation-Reduction
Reactions
20.1 The Meaning of Oxidation
and Reduction
20.2 Oxidation Numbers
20.3 Describing Redox
Equations
1
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20.3 Describing Redox Equations >
CHEMISTRY
& YOU
Why does cut fruit turn brown?
You have probably
noticed that the flesh
of an apple turns
brown after you
remove the skin. The
apple is still safe to
eat; it just doesn’t
look as appetizing.
2
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20.3 Describing Redox Equations > Identifying Redox
Reactions
Identifying Redox Reactions
What are the two classes of chemical
reactions?
3
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20.3 Describing Redox Equations > Identifying Redox
Reactions
All chemical reactions can be assigned to
one of two classes.
One class of chemical reactions is oxidationreduction (redox) reactions, in which
electrons are transferred from one reacting
species to another.
The other class includes all other reactions,
in which no electron transfer occurs.
4
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20.3 Describing Redox Equations > Identifying Redox
Reactions
Many single-replacement reactions,
combination reactions, decomposition
reactions, and combustion reactions are
redox reactions.
• Potassium metal
reacts violently
with water to
produce hydrogen
gas (which ignites)
and potassium
hydroxide.
5
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20.3 Describing Redox Equations > Identifying Redox
Reactions
Many single-replacement reactions,
combination reactions, decomposition
reactions, and combustion reactions are
redox reactions.
• Zinc metal reacts
vigorously with
hydrochloric acid
to produce
hydrogen gas and
zinc chloride.
6
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20.3 Describing Redox Equations > Identifying Redox
Reactions
Examples of reactions that are not redox
reactions include double-replacement
reactions and acid-base reactions.
7
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20.3 Describing Redox Equations > Identifying Redox
Reactions
During an electrical storm,
oxygen molecules and
nitrogen molecules in air
react to form nitrogen
monoxide.
N2(g) + O2(g) → 2NO(g)
How can you tell if this is a redox reaction?
8
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20.3 Describing Redox Equations > Identifying Redox
Reactions
During an electrical storm,
oxygen molecules and
nitrogen molecules in air
react to form nitrogen
monoxide.
N2(g) + O2(g) → 2NO(g)
How can you tell if this is a redox reaction?
• The oxidation number of nitrogen increases from 0 to +2.
• The oxidation number of oxygen decreases from 0 to –2.
• The reaction between nitrogen and oxygen to form
nitrogen monoxide is a redox reaction.
9
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20.3 Describing Redox Equations > Identifying Redox
Reactions
Many reactions in which color changes
occur are redox reactions.
• An example is shown below.
MnO4–(aq) + Br–(aq) → Mn2+(aq) + Br2(aq)
Permanganate Bromide ion
ion (purple)
(colorless)
10
Manganese(III)
ion (colorless)
Bromine
(brown)
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20.3 Describing Redox Equations >
CHEMISTRY
& YOU
Some fruits, including apples, turn
brown when you cut them. What do you
think is happening on the surface of the
fruit that causes it to turn brown?
11
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20.3 Describing Redox Equations >
CHEMISTRY
& YOU
Some fruits, including apples, turn
brown when you cut them. What do you
think is happening on the surface of the
fruit that causes it to turn brown?
Oxygen in air reacts with
chemicals on the surface of the
cut fruit. The oxygen oxidizes
the chemicals in the fruit,
causing a redox reaction and
therefore the color change.
12
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20.3 Describing Redox Equations >
Sample Problem 20.5
Identifying Redox Reactions
Use the change in oxidation number to
identify whether each reaction is a redox
reaction or a reaction of some other type. If
a reaction is a redox reaction, identify the
element reduced, the element oxidized, the
reducing agent, and the oxidizing agent.
a. Cl2(g) + 2NaBr(aq) → 2NaCl(aq) + Br2(aq)
b. 2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)
13
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20.3 Describing Redox Equations >
Sample Problem 20.5
1 Analyze Identify the relevant concepts.
• If changes in oxidation number occur,
the reaction is a redox reaction.
• The element whose oxidation number
increases is oxidized and is the reducing
agent.
• The element whose oxidation number
decreases is reduced and is the
oxidizing agent.
14
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20.3 Describing Redox Equations >
Sample Problem 20.5
2 Solve Apply concepts to this situation.
a. Assign oxidation numbers.
0
+1 –1
+1 –1
0
Cl2(g) + 2NaBr(aq) → 2NaCl(aq) + Br2(aq)
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20.3 Describing Redox Equations >
Sample Problem 20.5
2 Solve Apply concepts to this situation.
a. Interpret the change (or lack of change) in
oxidation numbers to identify if the reaction
is a redox reaction.
0
+1 –1
+1 –1
0
Cl2(g) + 2NaBr(aq) → 2NaCl(aq) + Br2(aq)
• This is a redox reaction.
• The chlorine is reduced.
• The bromide ion is oxidized.
• Chlorine is the oxidizing agent; the bromide ion is
the reducing agent.
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20.3 Describing Redox Equations >
Sample Problem 20.5
2 Solve Apply concepts to this situation.
b. Assign oxidation numbers.
+1 –2 +1
+1 +6 –2
+1 +6 –1
+1 –2
2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)
17
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20.3 Describing Redox Equations >
Sample Problem 20.5
2 Solve Apply concepts to this situation.
b. Interpret the change (or lack of change) in
oxidation numbers to identify if the reaction
is a redox reaction.
+1 –2 +1
+1 +6 –2
+1 +6 –1
+1 –2
2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)
• None of the elements change in
oxidation number.
• This is not a redox reaction.
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This is an
acid-base
(neutralization)
reaction.
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20.3 Describing Redox Equations >
Which of the following are redox
reactions?
A. NH3 + HCl → NH4Cl
B. SO3 + H2O → H2SO4
C. NaOH + HCl → NaCl + H2O
D. H2S + NHO3 → H2SO4 + NO2 + H2O
19
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20.3 Describing Redox Equations >
Which of the following are redox
reactions?
A. NH3 + HCl → NH4Cl
B. SO3 + H2O → H2SO4
C. NaOH + HCl → NaCl + H2O
D. H2S + NHO3 → H2SO4 + NO2 + H2O
20
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20.3 Describing Redox Equations > Balancing Redox
Equations
Balancing Redox Equations
What are two different methods for
balancing a redox equation?
21
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20.3 Describing Redox Equations > Balancing Redox
Equations
Two different methods for balancing
redox equations are the oxidationnumber-change method and the
half-reaction method.
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20.3 Describing Redox Equations > Balancing Redox
Equations
Two different methods for balancing
redox equations are the oxidationnumber-change method and the
half-reaction method.
• These two methods are based on the fact
that the total number of electrons gained
in reduction must equal the total number
of electrons lost in oxidation.
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Oxidation-Number Changes
In the oxidation-number-change method,
you balance a redox equation by comparing
the increases and decreases in oxidation
numbers.
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Oxidation-Number Changes
To use this method, start with the skeleton
equation for the redox reaction.
Fe2O3(s) + CO(g) → Fe(s) + CO2(g) (unbalanced)
In a blast furnace, air is blown
through a combination of iron
ore and coke. The carbon
monoxide produced from the
oxidation of coke reduces the
Fe3+ ions to metallic iron.
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Oxidation-Number Changes
Step 1: Assign oxidation numbers to all the
atoms in the equation.
• Write the numbers above the atoms.
• The oxidation number is stated per
atom.
+3 –2
+2 –2
0
+4 –2
Fe2O3(s) + CO(g) → Fe(s) + CO2(g)
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Oxidation-Number Changes
Step 2: Identify which atoms are oxidized and
which are reduced.
• Iron is reduced.
• Carbon is oxidized.
+3 –2
+2 –2
0
+4 –2
Fe2O3(s) + CO(g) → Fe(s) + CO2(g)
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Oxidation-Number Changes
Step 3: Use one bracketing line to connect the
atoms that undergo oxidation and
another such line to connect those
that undergo reduction.
• Write the oxidation-number change at
the midpoint of each line.
+2 (oxidation)
+3 –2
+2 –2
0
+4 –2
Fe2O3(s) + CO(g) → Fe(s) + CO2(g)
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–3 (reduction)
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Oxidation-Number Changes
Step 4: Make the total increase in oxidation
number equal to the total decrease in
oxidation number by using appropriate
coefficients.
• The oxidation-number increase should be
multiplied by 3 and the decrease by 2.
3 × (+2) = +6
+3 –2
+2 –2
0
+4 –2
Fe2O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g)
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2 × (–3) = –6
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Oxidation-Number Changes
Step 5: Finally, make sure the equation is
balanced for both atoms and charge.
• If necessary, finish balancing the
equation by inspection.
Fe2O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g)
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20.3 Describing Redox Equations >
Sample Problem 20.6
Balancing Redox Equations by
Oxidation-Number Change
Balance this redox equation by using the oxidationnumber-change method.
K2Cr2O7(aq) + H2O(l) + S(s) → KOH(aq) + Cr2O3(s) + SO2(g)
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20.3 Describing Redox Equations >
Sample Problem 20.6
1 Analyze Identify the relevant concepts.
You can balance redox equations by
determining changes in oxidation numbers
and applying the five steps.
32
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20.3 Describing Redox Equations >
Sample Problem 20.6
2 Solve Apply concepts to this situation.
Step 1: Assign oxidation numbers.
+1 +6 –2
+1 –2
0
+1–2+1
+3 –2
+4 –2
K2Cr2O7(aq) + H2O(l) + S(s) → KOH(aq) + Cr2O3(s) + SO2(g)
33
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20.3 Describing Redox Equations >
Sample Problem 20.6
2 Solve Apply concepts to this situation.
Step 2: Identify the atoms that are oxidized
and reduced.
+1 +6 –2
+1 –2
0
+1–2+1
+3 –2
+4 –2
K2Cr2O7(aq) + H2O(l) + S(s) → KOH(aq) + Cr2O3(s) + SO2(g)
• Cr is reduced.
• S is oxidized.
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20.3 Describing Redox Equations >
Sample Problem 20.6
2 Solve Apply concepts to this situation.
Step 3: Connect the atoms that change in
oxidation number. Indicate the signs
and magnitudes of the changes.
–3
+6
0
+3
+4
K2Cr2O7(aq) + H2O(l) + S(s) → KOH(aq) + Cr2O3(s) + SO2(g)
+4
35
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20.3 Describing Redox Equations >
Sample Problem 20.6
2 Solve Apply concepts to this situation.
Step 4: Balance the increase and decrease in
oxidation numbers.
(4)(–3) = –12
+6
0
+3
+4
2K2Cr2O7(aq) + H2O(l) + 3S(s) → KOH(aq) + 2Cr2O3(s) + 3SO2(g)
(3)(+4) = +12
Four chromium atoms must be reduced (4 × (–3) = –12 decrease) for
every three sulfur atoms that are oxidized (3 × (+4) = +12 increase).
Put the coefficient 3 in front of S and SO2, and the coefficient 2 in front
of K2Cr2O7 and Cr2O3.
36
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20.3 Describing Redox Equations >
Sample Problem 20.6
2 Solve Apply concepts to this situation.
Step 5: Check the equation and balance by
inspection if necessary.
2K2Cr2O7(aq) + 2H2O(l) + 3S(s) →
4KOH(aq) + 2Cr2O3(s) + 3SO2(g)
The coefficient 4 in front of KOH balances
potassium. The coefficient 2 in front of
H2O balances hydrogen and oxygen.
37
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Half-Reactions
A half-reaction is an equation showing just
the oxidation or just the reduction that takes
place in a redox reaction.
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Half-Reactions
In the half-reaction method, you write and
balance the oxidation and reduction halfreactions separately before combining them
into a balanced redox equation.
• The procedure is different, but the outcome is
the same as with the oxidation-number-change
method.
39
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Half-Reactions
Sulfur is an element that can
have several different oxidation
numbers.
• The oxidation of sulfur by nitric
acid in aqueous solution is one
example of a redox reaction that
can be balanced by following the
steps of the half-reaction method.
Oxidation Numbers of
Sulfur in Different
Substances
Substance
Oxidation
number
SO3
+6
SO2
+4
Na2S2O3
+2
S2Cl2
+1
S
0
H2S
–2
S(s) + HNO3(aq) → SO2(g) + NO(g) + H2O(l)
(unbalanced)
40
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Half-Reactions
Step 1: Write the unbalanced equation in ionic
form.
• Only HNO3 is ionized.
• The products are covalent compounds.
S(s) + H+(aq) + NO3–(aq) → SO2(g) + NO(g) + H2O(l)
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Half-Reactions
Step 2: Write separate half-reactions for the
oxidation and reduction processes.
• Sulfur is oxidized.
• Nitrogen is reduced.
• H+ ions and H2O are not included because
they are neither oxidized nor reduced.
0
+4
Oxidation half-reaction: S(s) → SO2(g)
+5
+2
Reduction half-reaction: NO3–(aq) → NO(g)
42
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Half-Reactions
Step 3: Balance the atoms in the halfreactions.
The half-reaction method is very useful
in balancing equations for reactions that
take place in acidic or basic solutions.
• In acidic solutions, H2O and H+(aq) can be
used to balance oxygen and hydrogen as
needed.
• In basic solution, H2O and OH– are used to
balance these species.
43
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Half-Reactions
Step 3: Balance the atoms in the halfreactions.
a. Balance the oxidation half-reaction.
• Sulfur is already balanced, but oxygen is not.
• This reaction takes place in acid solution, so
H2O and H+(aq) are present and can be used
to balance oxygen and hydrogen as needed.
2H2O(l) + S(s) → SO2(g) + 4H+(aq)
44
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Half-Reactions
Step 3: Balance the atoms in the halfreactions.
b. Balance the reduction half-reaction.
• Nitrogen is already balanced.
• Add two molecules of H2O on the right to
balance the oxygen.
• Four hydrogen ions must be added to the left
to balance hydrogen.
4H+(aq) + NO3–(aq) → NO(g) + 2H2O(l)
45
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Half-Reactions
Step 4: Add enough electrons to one side of
each half-reaction to balance the
charges.
• Neither half-reaction is balanced for
charge.
• Four electrons are needed on the right
side in the oxidation half-reaction.
Oxidation: 2H2O(l) + S(s) → SO2(g) + 4H+(aq) + 4e–
46
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Half-Reactions
Step 4: Add enough electrons to one side of
each half-reaction to balance the
charges.
• Neither half-reaction is balanced for
charge.
• Three electrons are needed on the left
side in the reduction half-reaction.
Reduction: 4H+(aq) + NO3–(aq) + 3e– → NO(g) + 2H2O(l)
47
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Half-Reactions
Step 5: Multiply each half-reaction by an
appropriate number to make the
numbers of electrons equal in both.
• The number of electrons lost in
oxidation must equal the number of
electrons gained in reduction.
Oxidation: 6H2O(l) + 3S(s) → 3SO2(g) + 12H+(aq) + 12e–
Reduction: 16H+(aq) + 4NO3–(aq) + 12e– →
4NO(g) + 8H2O(l)
48
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Half-Reactions
Step 6: Add the balanced half-reactions to
show an overall equation.
6H2O(l) + 3S(s) + 16H+(aq) + 4NO3–(aq) + 12e– →
3SO2(g) + 12H+(aq) + 12e– + 4NO(g) + 8H2O(l)
49
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Half-Reactions
Step 6: Add the balanced half-reactions to
show an overall equation.
• Then subtract terms that appear on
both sides of the equation.
3S(s) + 4H+(aq) + 4NO3–(aq) →
3SO2(g) + 4NO(g) + 2H2O(l)
50
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20.3 Describing Redox Equations > Balancing Redox
Equations
Using Half-Reactions
Step 7: Add the spectator ions and balance the
equation.
• Recall that spectator ions are present
but do not participate in or change
during a reaction.
• There are no spectator ions in this
particular example.
3S(s) + 4HNO3(aq) → 3SO2(g) + 4NO(g) + 2H2O(l)
51
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20.3 Describing Redox Equations >
Sample Problem 20.7
Balancing Redox Equations by
Half-Reactions
Balance this redox equation
using the half-reaction method.
KMnO4(aq) + HCl(l) → MnCl2(aq) + Cl2(g) + KCl(aq)
52
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20.3 Describing Redox Equations >
Sample Problem 20.7
1 Analyze Identify the relevant concepts.
You can use the seven steps of the
half-reaction method.
53
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20.3 Describing Redox Equations >
Sample Problem 20.7
2 Solve Apply concepts to this problem.
Step 1:Write the equation in ionic form.
K+(aq) + MnO4–(aq) + H+(aq) + Cl–(aq) →
Mn2+(aq) + 2Cl–(aq) + Cl2(g) + H2O(l) + K+(aq) + Cl–(aq)
54
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20.3 Describing Redox Equations >
Sample Problem 20.7
2 Solve Apply concepts to this problem.
Step 2: Write half-reactions. Determine
the oxidation and reduction
process.
–1
0
Oxidation half-reaction: Cl– → Cl2
+7
+2
Reduction half-reaction: MnO4– → Mn2+
55
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20.3 Describing Redox Equations >
Sample Problem 20.7
2 Solve Apply concepts to this problem.
Step 3: Balance the atoms in each halfreaction.
• The solution is acidic, so use H2O and H+ to
balance the oxygen and hydrogen.
Oxidation: 2Cl–(aq) → Cl2(g) (atoms balanced)
Reduction: MnO4–(aq) + 8H+(aq) →
Mn2+(aq) + 4H2O(l) (atoms balanced)
56
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20.3 Describing Redox Equations >
Sample Problem 20.7
2 Solve Apply concepts to this problem.
Step 4: Balance the charges by adding
electrons.
Oxidation: 2Cl–(aq) → Cl2(g) + 2e– (charges balanced)
Reduction: MnO4–(aq) + 8H+(aq) + 5e– →
Mn2+(aq) + 4H2O(l) (charges balanced)
57
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20.3 Describing Redox Equations >
Sample Problem 20.7
2 Solve Apply concepts to this problem.
Step 5: Make the numbers of electrons
equal.
• Multiply the oxidation half-reaction by 5
and the reduction half-reaction by 2.
Oxidation: 10Cl–(aq) → 5Cl2(g) + 10e–
Reduction: 2MnO4–(aq) + 16H+(aq) + 10e– →
2Mn2+(aq) + 8H2O(l)
58
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20.3 Describing Redox Equations >
Sample Problem 20.7
2 Solve Apply concepts to this problem.
Step 6: Add the half-reactions. Then,
subtract the terms that appear on
both sides.
10Cl–(aq) + 2MnO4–(aq) + 16H+(aq) + 10e– →
5Cl2(g) + 10e– + 2Mn2+(aq) + 8H2O(l)
59
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20.3 Describing Redox Equations >
Sample Problem 20.7
2 Solve Apply concepts to this problem.
Step 7: Add the spectator ions, making
sure the charges and atoms are
balanced.
10Cl– + 2MnO4– + 2K+ + 16H+ + 6Cl– →
5Cl2 + 2Mn2+ + 4Cl– + 8H2O + 2K+ + 2Cl–
60
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20.3 Describing Redox Equations >
Sample Problem 20.7
2 Solve Apply concepts to this problem.
Combine the spectator and nonspectator
Cl– on each side.
16Cl–(aq) + 2MnO4–(aq) + 2K+(aq) + 16H+(aq) →
5Cl2(g) + 2Mn2+(aq) + 6Cl–(aq) + 8H2O(l) + 2K+(aq)
61
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20.3 Describing Redox Equations >
Sample Problem 20.7
2 Solve Apply concepts to this problem.
Show the balanced equation for the
substances given in the question (rather
than for ions).
2KMnO4(aq) + 16HCl(aq) →
2MnCl2(aq) + 5Cl2(g) + 8H2O(l) + 2KCl(aq)
62
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20.3 Describing Redox Equations >
Use the half-reaction method to
balance the following redox equation.
FeCl3 + H2S → FeCl2 + HCl + S
63
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20.3 Describing Redox Equations >
Use the half-reaction method to
balance the following redox equation.
FeCl3 + H2S → FeCl2 + HCl + S
Oxidation: H2S → 2H+ + S + 2e–
Reduction: 2Fe3+ + 2e– → 2Fe2+
2FeCl3 + H2S → 2FeCl2 + 2HCl + S
64
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20.3 Describing Redox Equations > Key Concepts
One class of chemical reactions is
oxidation reduction (redox) reactions, in
which electrons are transferred from one
reacting species to another. The other
class includes all other reactions, in which
no electron transfer occurs.
To balance a redox equation using the
oxidation-number-change method, the total
increase in oxidation number of the
species oxidized must be balanced by the
total decrease in the oxidation number of
the species reduced.
65
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20.3 Describing Redox Equations > Key Concepts
To balance a redox reaction using halfreactions, write separate half-reactions for
the oxidation and the reduction. After you
balance atoms in each half-reaction,
balance electrons gained in the reduction
with electrons lost in the oxidation.
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20.3 Describing Redox Equations > Glossary Terms
• oxidation-number-change method: a method
of balancing a redox equation by comparing the
increases and decreases in oxidation numbers
• half-reaction: an equation showing either the
oxidation or the reduction that takes place in a
redox reaction
• half-reaction method: a method of balancing a
redox equation by balancing the oxidation and
reduction half-reactions separately before
combining them into a balanced redox equation
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20.3 Describing Redox Equations >
BIG IDEA
Reactions
Redox equations can be balanced by two
methods, the oxidation-number-change
method and balancing the oxidation and
reduction half-reactions.
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20.3 Describing Redox Equations >
END OF 20.3
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