Download Chapter 18

Chapter 18
Oxidation–Reduction
Reactions and
Electrochemistry
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Chapter 18
Table of Contents
• Reaction rates decrease with time because reactant
concentrations decrease as reactants are converted to
products.
• Less concentration lead fewer collision，so the rate of
reaction decrease。
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Chapter 18
Table of Contents
（1）When the temperature is increased, the average velocity of the
particles is increased. The result is that the particles will collide more
frequently, because the particles move around faster and will
encounter more reactant particles.
（2）The major effect of increasing the temperature is that more of the
particles that collide will have the amount of energy needed to have an
effective collision. In other words, more particles with higher average
kinetic energy will have the necessary activation energy.
Increasing the temperature increases the rate of reaction because
the particles collide more often and with more energy.
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Chapter 18
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CO2 + H2 ⇔ CO + H2O
I
C
E
0.00
0.00
+X
+X
X
X
[𝐶𝑂][𝐻2𝑂]
K=
=0.279
[𝐶𝑂2][𝐻2]
[4.00−𝑥]
[𝑥]
=0.528
[CO2]=[H2]=2.62
4.00 4.00
-X
-X
4.00-X 4.00-X
[4.00−𝑥]2
=
[𝑥]2
x=2.62
[CO]=[H2O]=1.38
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Chapter 18
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Chapter 18
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Chapter 18
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Chapter 18
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The apple is oxidized.
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Chapter 18
Table of Contents
Oxidation–reduction reactions
Among the first reactions studied by early scientists were
those that involved oxygen.
• The combustion of fuels and the reactions of metals with
oxygen to give oxides were described by the word
oxidation.
• The removal of oxygen from metal oxides to give the
metals in their elemental forms was described as
reduction.
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Chapter 18
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3𝐹𝑒 + 2 𝑂2 → 𝐹𝑒3𝑂4
The combustion of fuels and the
reactions of metals with oxygen to give
oxides were described by the word
oxidation.
2𝐶𝑢0 + 𝐶 → 2𝐶𝑢 + 𝐶𝑂2
The removal of oxygen from metal
oxides to give the metals in their
elemental forms was described as
reduction.
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Chapter 18
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Summary
2𝐶𝑢0 + 𝐶 → 2𝐶𝑢 + 𝐶𝑂2
Cu is lose oxygen , reduction occur ,Cu is reduced.
C is gain oxygen , oxidation occur ,C is oxidized
Oxidation–reduction reactions
+2→0，gain electron ———— reduction
0 →+4，lose electron ———— oxidation
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Section 18.1
Oxidation–Reduction Equations
Metal–Nonmetal Oxidation–Reduction Reactions
• Oxidation–reduction reactions: Chemical reactions
involving the transfer of electrons
– Also known as redox reactions
– Oxidation: Loss of electrons
– Reduction: Gain of electrons
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Section 18.3
Oxidation–Reduction Reactions Between Nonmetals
Oxidising and Reducing Agents
• 2Na(s) + Cl2(g)  2NaCl(s)
• Na  Oxidized
– Lose Electron
– Na is also called the reducing agent
• Cl2  Reduced
– Gain Electron
−Cl2 is also called the oxidizing agent
oil
rig
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Section 18.3
Oxidation–Reduction Reactions Between Nonmetals
Oxidising and Reducing Agents (continued)
• CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
• C  Oxidized
– CH4 is the reducing agent
• O2  Reduced
– O2 is the oxidizing agent
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Section 18.1
Oxidation–Reduction Equations
Exercise, 1
What happens to Sn(II) in the given reaction?
Sn2+ + 2Fe3+ → Sn4+ + 2Fe2+
a) It gains electrons
b) It is reduced
c) It is oxidized
d) It is neither oxidized nor reduced
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Chapter 18
Table of Contents
Patterns can be etched onto aluminum with a beam of chlorine
atoms. Identify the substances oxidized and reduced, and the
oxidizing and reducing agents in the reaction of aluminum and
chlorine to form aluminum chloride.
2𝐴𝑙 + 3𝐶𝑙2 → 2𝐴𝑙𝐶𝑙3
Aluminum is oxidized and is, therefore, the reducing agent.
Chlorine is reduced and is, therefore, the oxidizing agent.
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Section 18.2
Oxidation States
Oxidation States
• Helps keep track of electrons in oxidation–reduction
reactions by assigning charges to the various atoms in a
compound
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Section 18.2
Oxidation States
Rules for Assigning Oxidation States
• Oxidation state of :
– (1)All the atoms of free elements have oxidation numbers of zero.
– (2)Metals in Groups 1A, 2A, and Al have +1, +2, and +3 oxidation
numbers, respectively.
– (3)H and F, in compounds, have +1 and −1 oxidation numbers,
respectively.
– (4)Hydrogen is +1 in covalent compounds with nonmetals.Except
when bonded to Group I or Group II, when it forms hydrides, -1.
– (5)Oxygen has a −2 oxidation number. peroxides (compounds
containing the O2 -2 group), in which each oxygen is assigned an
oxidation state of -1.
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Section 18.2
Oxidation States
Rules for Assigning Oxidation States (continued)
•
•
•
•
(6)Group 7A elements have a −1 oxidation number.
(7)Group 6A elements have a −2 oxidation number.
(8)Group 5A elements have a −3 oxidation number.
(9)In binary compounds, the most electronegative
element is assigned a negative oxidation state.
19
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Section 18.2
Oxidation States
Exercise, 2
Find the oxidation states for each of the
elements in the following compounds:
•
•
•
•
•
K2Cr2O7
CO32−
MnO2
PCl5
SF4
K = +1; Cr = +6; O = −2
C = +4; O = −2
Mn = +4; O = −2
P = +5; Cl = −1
S = +4; F = −1
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20
Section 18.3
Oxidation–Reduction Reactions Between Nonmetals
Redox Characteristics
•
•
•
Transfer of electrons
Transfer may occur to form ions
Oxidation: Increase in oxidation state
–
–
•
Loss of electrons
One that gets oxidized is called the reducing
agent
Reduction: Decrease in oxidation state
–
–
Gain of electrons
One that gets reduced is called the oxidizing
agent
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Section 18.3
Oxidation–Reduction Reactions Between Nonmetals
Concept Check
Which of the following are oxidation–reduction reactions?
Identify the oxidizing agent and the reducing agent.
a)Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
b)Cr2O72−(aq) + 2OH−(aq)  2CrO42−(aq) + H2O(l)
c)2CuCl(aq)  CuCl2(aq) + Cu(s)
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Section 18.3
Oxidation–Reduction Reactions Between Nonmetals
Goals for Balancing Chemical Equations
1. The number of atoms of each element on both sides
of the equation is the same and therefore mass is
conserved.
2. The sum of the positive and negative charges is the
same on both sides of the equation and therefore charge
is conserved.
3. The number of losed electrons and gained electrons is
same
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Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Half-Reactions
•
•
Overall reaction is split into two half-reactions, one
involving oxidation and one reduction
Have electrons as reactants or products
8H+ + MnO4− + 5Fe2+ → Mn2+ + 5Fe3+ + 4H2O
Reduction: 8H+ + MnO4− + 5e− → Mn2+ + 4H2O
Oxidation: 5Fe2+ → 5Fe3+ + 5e−
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Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Half-Reaction Method for Balancing Equations for Oxidation–
Reduction Reactions Occurring in Acidic Solution
• Identify and write the equations for the oxidation
and reduction half-reactions
• For each half-reaction:
a)
b)
c)
d)
Balance all the elements except H and O
Balance O using H2O
Balance H using H+
Balance the charge using electrons
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Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Half-Reaction Method for Balancing Equations for Oxidation–
Reduction Reactions Occurring in Acidic Solution (continued 1)
• If necessary, multiply one or both balanced halfreactions by an integer to equalize the number of
electrons transferred in the two half-reactions
• Add the half-reactions, and cancel identical
species that appear on both sides
• Check that the elements and charges are
balanced
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Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Half-Reaction Method for Balancing Equations for Oxidation–
Reduction Reactions Occurring in Acidic Solution (continued 2)
Cr2O72−(aq) + SO32−(aq)  Cr3+(aq) + SO42−(aq)
• How should this equation be balanced?
• Steps:
– Separate into half-reactions
– Balance elements except H and O
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Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Mechanism of Half-Reactions
Cr2O72−(aq)  2Cr3+(aq)
SO32−(aq)  SO42−(aq)
•
How many electrons are involved in each halfreaction?
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28
Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Mechanism of Half-Reactions (continued 1)
6e− + Cr2O72−(aq)  2Cr3+(aq)
SO32−(aq)  SO42−(aq) + 2e−
• How should the oxygen atoms be balanced?
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29
Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Mechanism of Half-Reactions (continued 2)
6e− + Cr2O72−(aq)  Cr3+(aq) + 7H2O
H2O + SO32−(aq)  SO42−(aq) + 2e−
• How should the hydrogen atoms be balanced?
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30
Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Mechanism of Half-Reactions (continued 3)
•
This reaction occurs in an acidic solution
14H+ + 6e− + Cr2O72−  2Cr3+ + 7H2O
H2O + SO32−  SO42− + 2e− + 2H+
• How should the electrons be balanced?
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Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Mechanism of Half-Reactions (continued 4)
14H+ + 6e− + Cr2O72−  2Cr3+ + 7H2O
3[H2O + SO32−  SO42− + 2e− + 2H+]
• Final balanced equation:
Cr2O72− + 3SO32− + 8H+  2Cr3+ + 3SO42− + 4H2O
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Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Exercise, 3
When the reaction Ce2+ + Co2+ → Ce3+ + Co is
balanced, the coefficient of Ce2+ is
a) 0
b) 1
c) 2
d) 3
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33
Section 18.4
Balancing Oxidation–Reduction Reactions by the Half-Reaction Method
Exercise, 4
Balance the following oxidation–reduction reaction
that occurs in acidic solution
Br−(aq) + MnO4−(aq)  Br2(l)+ Mn2+(aq)
10Br  (aq )  16H (aq ) + 2MnO 4  (aq )  5Br2 ( l ) + 2Mn2 (aq )  8H2O(l )
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34
Section 18.5
Electrochemistry: An Introduction
Electrochemistry
• Study of the interchange of chemical and electrical energy
• Two types of processes
– Production of an electric current from a chemical reaction
– Use of an electric current to produce a chemical change
Galvanic cells in which spontaneous oxidation-reduction
reactions produce electrical energy.
voltaic cell.
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35
Section 18.5
Electrochemistry: An Introduction
Making an Electrochemical Cell
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Section 18.5
Electrochemistry: An Introduction
根据科学家之间的协议，阳极和阴极的名称是根据电极上发生
的反应的性质来分配的。
如果反应是氧化，则电极被称为阳极；
如果是还原，电极就叫做阴极。
If the reaction is oxidation,
the electrode is called the anode;
If the reaction is reduction,
the electrode is called the cathode.
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Section 18.5
Electrochemistry: An Introduction
Making an Electrochemical Cell (continued 1)
• Electrodes——An electrode is strip of metal on which the
reaction takes place.
——Electrodes can be defined as metallic conduction, that
are used to make electrical contact with a non-metallic part
of the circuit.
1. The electrode is composed of
metals with different activity, in
which the active metal that reacts
with the electrolyte solution acts as
the negative electrode
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38
Section 18.5
Electrochemistry: An Introduction
• Sometimes, both the oxidized and reduced forms of the
reactants in a half-cell are soluble and neither can be
used as an electrode. In these cases, an inert electrode
composed of platinum, graphite, or gold is used to
provide a site for electron transfer.
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39
Section 18.5
Electrochemistry: An Introduction
Electrochemical Battery (Galvanic Cell)
• Device powered by an oxidation–reduction reaction
– Oxidizing agent is separated from the reducing agent so
that the electrons travel through a wire from the reducing
agent to the oxidizing agent
• Anode: Electrode where oxidation occurs
• Cathode: Electrode where reduction occurs
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40
Section 18.5
Electrochemistry: An Introduction
Making an Electrochemical Cell (continued 1)
• If electrons flow through
the wire, charge builds up
– To balance the charge in
each compartment, the
solutions must be
connected so that ions
can flow
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41
Section 18.5
Electrochemistry: An Introduction
Making an Electrochemical Cell (continued 1)
• Current produced in the
wire by this electron flow
can be directed through a
device to do useful work
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42
Section 18.5
Electrochemistry: An Introduction
Making an Electrochemical Cell
• Salt bridges or porous disks can be used to
connect the half-cells
– Allow ion flow and thus complete the circuit
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43
Section 18.5
Electrochemistry: An Introduction
Making an Electrochemical Cell
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44
Electrons are responsible for the current flow.
In the solution, conduction of electricity is caused by migration of
ions
Electrochemistry
© 2015 Pearson Education, Inc.
Section 18.5
Electrochemistry: An Introduction
Electrolysis
• Process where electrical energy is used to produce a
chemical change
• Example- Current can be passed through water to
produce hydrogen and oxygen
Electrical energy
2H2O(l ) 
 2H2 ( g ) + O2 ( g )
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46
Section 18.6
Batteries
Lead Storage Battery
• Anode reaction: Oxidation
Pb + H2SO4  PbSO4 + 2H+ + 2e
• Cathode reaction: Reduction
PbO2 + H2SO4 + 2e + 2H+  PbSO4 + 2H2O
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47
Section 18.6
Batteries
Lead Storage Battery Overall Reaction
Pb(s) + PbO2(s) + 2H2SO4(aq)  2PbSO4(s) + 2H2O(l)
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48
Section 18.6
Batteries
Electric Potential
• “Pressure” on electrons to flow from one electrode to
another in a battery
• Measured in volts
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49
Section 18.6
Batteries
Dry Cell Batteries
• Do not contain a liquid electrolyte
– Acid version
•
•
Anode reaction: Oxidation
Zn  Zn2+ + 2e
Cathode reaction: Reduction
2NH4+ + 2MnO2 + 2e  Mn2O3 + 2NH3 + H2O
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50
Section 18.6
Batteries
Dry Cell Batteries (continued 1)
– Alkaline version
• Anode reaction: Oxidation
•
Zn + 2OH  ZnO + H2O + 2e
Cathode reaction: Reduction
2MnO2 + H2O + 2e  Mn2O3 + 2OH
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51
Section 18.6
Batteries
Dry Cell Batteries (continued 2)
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52
Section 18.6
Batteries
Dry Cell Batteries (continued 3)
– Other Types
• Silver cell: Zn anode, Ag2O cathode
• Mercury cell: Zn anode, HgO cathode
• Lithium ion battery: Anode is a porous form of graphite
(C) into which Li+ ions have been inserted, and cathode
is a metal oxide such as LiCoO2
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53
Section 18.7
Corrosion
Corrosion
• Oxidation of metals to form mainly oxides and sulfides
• Prevented through the application of a coating
– Paint
– Metal Plating
– Alloying
• Some metals develop an oxide coating, which protects
their internal atoms against further oxidation
– Examples- Aluminum, chromium, nickel, and tin
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54
Section 18.7
Corrosion
Cathodic Protection
• Employed to protect steel in
buried pipelines and fuel tanks
• Metal that furnishes electrons
easily than iron is connected
by a wire to the pipeline
– Magnesium is used because it
is a better reducing agent than
iron
• Must be replaced periodically
because it dissolves when
oxidation occurs
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55
Section 18.8
Electrolysis
Electrolysis (continued)
• Forcing a current through a cell to produce a chemical
change that would not otherwise occur
• Electrolysis of water to produce hydrogen and oxygen
occurs whenever a current is forced through an aqueous
solution
forced electric current
2H2O(l ) 
 2H2 ( g ) + O2 ( g )
• Used in the production of metals from their ores
– Metal produced in the greatest quantities by electrolysis is
aluminium
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56