basic principles of chemistry online southeast missouri state university cape girardeau, mo...
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Basic Principles of Chemistry OnlineSoutheast Missouri State University
Cape Girardeau, MO
Introductory Chemistry, 3rd EditionNivaldo Tro
Chapter 16Oxidation and
Reduction
2009, Prentice Hall
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Oxidation–Reduction Reactions• Oxidation–reduction reactions are also called redox
reactions.• All redox reactions involve the transfer of electrons
from one atom to another.• Spontaneous redox reactions are generally
exothermic, and we can use their released energy as a source of energy for other applications.Convert the heat of combustion into mechanical energy to
move our cars.Use electrical energy in a car battery to start our car engine.
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Combustion Reactions• Combustion reactions are always exothermic.
• In combustion reactions, O2 combines with all the elements in another reactant to make the products.
4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) + energy
CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) + energy
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Reverse of Combustion Reactions• Since combustion reactions are exothermic, their
reverse reactions are endothermic.
• The reverse of a combustion reaction involves the production of O2.
energy + 2 Fe2O3(s) → 4 Fe(s) + 3 O2(g)
energy + CO2(g) + 2 H2O(g) → CH4(g) + 2 O2(g)
• Reactions in which O2 is gained or lost are redox reactions.
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Oxidation and Reduction:One Definition
• When an element attaches to an oxygen during the course of a reaction it is generally being oxidized. In CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g), C is being
oxidized in this reaction, but H is not.
• When an element loses an attachment to oxygen during the course of a reaction, it is generally being reduced. In 2 Fe2O3(s) → 4 Fe(s) + 3 O2(g), the Fe is being reduced.
• One definition of redox is the gain or loss of O, but it is not the best.
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Another Oxidation–Reduction
• Consider the following reactions:
4 Na(s) + O2(g) → 2 Na2O(s)
2 Na(s) + Cl2(g) → 2 NaCl(s)
• The reaction involves a metal reacting with a nonmetal.• In addition, both reactions involve the conversion of
free elements into ions.
4 Na(s) + O2(g) → 2 Na+2O–(s)
2 Na(s) + Cl2(g) → 2 Na+Cl–(s)
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Oxidation and Reduction:Another Definition
• In order to convert a free element into an ion, the atoms must gain or lose electrons.Of course, if one atom loses electrons, another must accept
them.
• Reactions where electrons are transferred from one atom to another are redox reactions.
• Atoms that lose electrons are being oxidized, atoms that gain electrons are being reduced.
2 Na(s) + Cl2(g) → 2 Na+Cl–(s)Na → Na+ + 1 e– (oxidation)Cl2 + 2 e– → 2 Cl– (reduction)
Leo
Ger
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Practice—Identify the Element Being Oxidized and the Element Being
Reduced.
• 2 C(s) + O2(g) → 2 CO(g)
• Mg(s) + Cl2(g) → MgCl2(s)
• Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s)
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Practice—Identify the Element Being Oxidized and the Element Being
Reduced, Continued.
• 2 C(s) + O2(g) → 2 CO(g)
• Mg(s) + Cl2(g) → MgCl2(s)
• Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s)
C is oxidized because it is gaining an attachment to O. O is reduced; there has to be reduction and it’s the only other element.
Mg is oxidized because it is becoming a cation by losing electrons. Cl is reduced because it is becoming an anion by gaining electrons.
0 0 2+ −
Mg is oxidized because it is becoming a cation by losing electrons. Fe2+ is reduced because it is gaining electrons to become neutral.
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Oxidation–Reduction• Oxidation and reduction must occur simultaneously.
If an atom loses electrons, another atom must take them.
• The reactant that reduces an element in another reactant is called the reducing agent.The reducing agent contains the element that is oxidized.
• The reactant that oxidizes an element in another reactant is called the oxidizing agent.The oxidizing agent contains the element that is reduced.
2 Na(s) + Cl2(g) → 2 Na+Cl–(s)Na is oxidized, Cl is reduced.
Na is the reducing agent, Cl2 is the oxidizing agent.
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Practice—Identify the Oxidizing and Reducing Agents.
• 2 C(s) + O2(g) → 2 CO(g)
• Mg(s) + Cl2(g) → MgCl2(s)
• Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s)
C is oxidized because it is gaining attachment to O. O is reduced; there has to be reduction and it’s the only other element.
Mg is oxidized because it is becoming a cation by losing electrons. Cl is reduced because it is becoming an anion by gaining electrons.
0 0 2+ −
Mg is oxidized because it is becoming a cation by losing electrons. Fe2+ is reduced because it is gaining electrons to become neutral.
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Practice—Identify the Oxidizing and Reducing Agents, Continued.
• 2 C(s) + O2(g) → 2 CO(g)
• Mg(s) + Cl2(g) → MgCl2(s)
• Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s)
C is the reducing agent because it contains the element that is oxidized. O is the oxidizing agent because it contains the element that is reduced.
0 0 2+ −
Mg is the reducing agent because it contains the element that is oxidized. Cl2 is the oxidizing agent because it contains the element that is reduced.
Mg is the reducing agent because it contains the element that is oxidized. Fe2+ is the oxidizing agent because it contains the element that is reduced.
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Electron Bookkeeping• For reactions that are not metal + nonmetal, or do
not involve O2, we need a method for determining how the electrons are transferred.
• Chemists assign a number to each element in a reaction called an oxidation state that allows them to determine the electron flow in the reaction.Although they look like them, oxidation states are not
ion charges!Oxidation states are imaginary charges assigned based on
a set of rules. Ion charges are real, measurable charges.
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Rules for Assigning Oxidation States
• Rules are in order of priority.
1. Free elements have an oxidation state = 0. Na(s) = 0 and Cl2(g) = 0 in 2 Na(s) + Cl2(g)
2 NaCl(s).
2. Monoatomic ions have an oxidation state equal to their charge.
Na = +1 and Cl = -1 in NaCl(s).
3. a. The sum of the oxidation states of all the atoms or ions in a compound is 0.
Na = +1 and Cl = -1 in NaCl, and (+1) + (-1) = 0.
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Rules for Assigning Oxidation States, Continued
3. b. The sum of the oxidation states of all the atoms in a polyatomic ion equals the charge on the ion.
N = +5 and O = -2 in NO3–, (+5) + 3(-2) = -1.
4. a. Group I metals have an oxidation state of +1 in all their compounds.
Na = +1 in NaCl.
b. Group II metals have an oxidation state of +2 in all their compounds.
Mg = +2 in MgCl2.
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Rules for Assigning Oxidation States, Continued
5. In their compounds, nonmetals have oxidation states according to the table below.
Nonmetals higher on the table take priority.Nonmetal Oxidation state Example
F -1 CF4
H +1 CH4
O -2 CO2
Group 7A -1 CCl4
Group 6A -2 CS2
Group 5A -3 NH3
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Practice—Assign an Oxidation State to Each Element in the Following:
• F2
• Mg2+
• KCl
• SO2
• PO43−
• BaO2
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Practice—Assign an Oxidation State to Each Element in the Following,
Continued:
• F2 F = 0 (Rule 1)
• Mg2+ Mg = +2 (Rule 2)
• KCl K = +1 (Rule 4a) and Cl = -1 (Rule 5)
• SO2 O = -2 (Rule 5) and S = +4 (Rule 3a)
• PO43− O = -2 (Rule 5) and P = +5 (Rule 3b)
• BaO2 Ba = +2 (Rule 4b) and O = -1 (Rule 3a)
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Oxidation and Reduction:A Better Definition
• Oxidation occurs when an atom’s oxidation state increases during a reaction.
• Reduction occurs when an atom’s oxidation state decreases during a reaction.
CH4 + 2 O2 → CO2 + 2 H2O-4 +1 0 +4 –2 +1 -2
oxidationreduction
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Practice—Assign Oxidation States and Identify the Oxidizing and Reducing Agents
in Each of the Following:
• 3 H2S + 2 NO3– + 2 H+ S + 2 NO + 4 H2O
• MnO2 + 4 HBr MnBr2 + Br2 + 2 H2O
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• 3 H2S + 2 NO3– + 2 H+ S + 2 NO + 4 H2O
• MnO2 + 4 HBr MnBr2 + Br2 + 2 H2O
+1 -2 +5 -2 +1 0 +2 -2 +1 -2
oxidizing agentreducing agent
+4 -2 +1 -1 +2 -1 0 +1 -2
oxidationreduction
oxidation
reduction
reducing agentOxidizing agent
Practice—Assign Oxidation States and Identify the Oxidizing and Reducing Agents
in Each of the Following, Continued:
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Will a Reaction Take Place?
• Reactions that are energetically favorable are said to be spontaneous.They can happen, but the activation energy may
be so large that the rate is very slow.
• The relative reactivity of metals can be used to determine if some redox reactions are spontaneous.
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Single Displacement Reactions• Also known as single replacement reactions.• A more active free element displaces a less active
element in a compound.Metals displace metals or H.
Cu + 2 AgNO3 Cu(NO3)2 + 2 Ag
Mg + 2 HCl MgCl2 + H2
Nonmetals displace nonmetals.
2 KI + Br2 2 KBr + I2
Carbon displaces metals from oxides.
3 C + Fe2O3 3 CO + 2 Fe
• Always redox.
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Tendency to Lose Electrons• Some metals have a greater tendency to lose
electrons than others.Metallic-free elements are always oxidized.The greater the tendency of a metal to lose electrons,
the easier it is to oxidize.The greater the tendency of a metal to lose electrons,
the harder it is to reduce its cations.
• If Metal A has a greater tendency to lose electrons than Metal B, then:
A(s) + B+(aq) A+(aq) + B(s),but: A+(aq) + B(s) no reaction.
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KBaSrCaNaMgAlMnZnCrFeCdCoNiSnPbHSbAsBiCuHgAgPdPtAu
displace H2
fromcoldH2O
fromsteam
fromacids
reac
t wit
h O
2 in
the
air
to m
ake
oxid
es
Fe is above Cu, so Cu metalwill not displace Fe2+
KBaSrCaNaMgAlMnZnCrFeCdCoNiSnPbHSbAsBiCuHgAgPdPtAu
displace H2
fromcoldH2O
fromsteam
fromacids
reac
t wit
h O
2 in
the
air
to m
ake
oxid
es
Gold is at thebottom, so it isvery unreactive.
KBaSrCaNaMgAlMnZnCrFeCdCoNiSnPbHSbAsBiCuHgAgPdPtAu
displace H2
fromcoldH2O
fromsteam
fromacids
reac
t wit
h O
2 in
the
air
to m
ake
oxid
es
Zn is above H,so Zn will react with acids
Zn + Fe2+ Fe + Zn2+
Activity Series of Metals• Listing of metals by
reactivity.• Free metal higher on the
list displaces metal cation lower on the list.
• Metals above H will dissolve in acid:
Cu + Fe2+ no reactionZn + 2 H+ H2 + Zn2+
KBaSrCaNaMgAlMnZnCrFeCdCoNiSnPbHSbAsBiCuHgAgPdPtAu
displace H2
fromcoldH2O
fromsteam
fromacids
reac
t wit
h O
2 in
the
air
to m
ake
oxid
es
Fe is below Zn, so Zn metalwill displace Fe2+.
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Mg is aboveCu on theactivity series.Mg will react with
Cu2+ to form Mg2+
and Cu metal.
Cu will not react with Mg2+.
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Table of Oxidation Half-Reactions
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Table of Oxidation Half-Reactions, Continued
• Any oxidation half-reaction that is higher on the list will give a spontaneous reaction when combined with the reverse of a half-reaction that is lower on the list.The reverse of an oxidation half-reaction is a
reduction half-reaction.
• Metals will dissolve in acid if their oxidation half-reaction is above H2 2H++ 2e−.
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Electrical Current• When we talk about the current
of a liquid in a stream, we are discussing the amount of water that passes by in a given period of time.
• When we discuss electric current, we are discussing the amount of electric charge that passes a point in a given period of time.Whether as electrons flowing
through a wire or ions flowing through a solution.
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Redox Reactions and Current• Redox reactions involve the transfer of
electrons from one substance to another.• Therefore, redox reactions have the
potential to generate an electric current.• In order to use that current, we need to
separate the place where oxidation is occurring from the place that reduction is occurring.
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Electric Current Flowing Directly Between Atoms
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Electric Current Flowing Indirectly Between Atoms
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Electrochemical Cells• Electrochemistry is the study of redox reactions that
produce or require an electric current.• The conversion between chemical energy and electrical
energy is carried out in an electrochemical cell.• Spontaneous redox reactions take place in a voltaic cell.
Also known as galvanic cells.Batteries are voltaic cells.
• Nonspontaneous redox reactions can be made to occur in an electrolytic cell by the addition of electrical energy.
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Electrochemical Cells, Continued• Oxidation and reduction reactions kept separate.
Half-cells.
• Electron flow through a wire, along with ion flow through a solution, constitutes an electric circuit.
• Requires a conductive solid (metal or graphite) electrode to allow the transfer of electrons.Through external circuit.
• Ion exchange between the two halves of the system.Electrolyte.
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Electrodes• Anode
Electrode where oxidation occurs.Anions attracted to it.Connected to positive end of battery in
electrolytic cell.Loses weight in electrolytic cell.
• CathodeElectrode where reduction occurs.Cations attracted to it.Connected to negative end of battery in
electrolytic cell.Gains weight in electrolytic cell.
Electrode where plating takes place in electroplating.
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Voltaic Cell
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Current and Voltage
• The number of electrons that flow through the system per second is the current.Electrode surface area dictates the number of
electrons that can flow.
• The amount of force pushing the electrons through the wire is the voltage.The farther the metals are separated on the
activity series, the larger the voltage will be.
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Current
The amount of water that passesa point each secondis called the currentof the river.
The number of electrons that passa point each secondis called the currentof the electricity.
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Voltage
Gravity is the force pullingthe water downthe river.
Voltage is the force pushingthe electrons down the wire.
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Dead Battery
As the reactionproceeds, thereactants getconsumed andthe voltaic cell“dies.” The current decreasesuntil electronscan no longerflow throughthe wire.
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LeClanché’s Acidic Dry Cell• Electrolyte in paste form.
ZnCl2 + NH4Cl.Or MgBr2.
• Anode = Zn (or Mg).Zn(s) Zn2+(aq) + 2 e-
• Cathode = graphite rod.• MnO2 is reduced.
2 MnO2(s) + 2 NH4+(aq) + 2 H2O(l) + 2 e- 2 NH4OH(aq) + 2 Mn(O)OH(s)
• Cell voltage = 1.5 v.• Expensive, nonrechargeable, heavy, easily
corroded.
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Alkaline Dry Cell• Same basic cell as acidic dry cell, except
electrolyte is alkaline KOH paste.
• Anode = Zn (or Mg).Zn(s) Zn2+(aq) + 2 e-
• Cathode = brass rod.
• MnO2 is reduced.2 MnO2(s) + 2 NH4
+(aq) + 2 H2O(l) + 2 e- 2 NH4OH(aq) + 2 Mn(O)OH(s)
• Cell voltage = 1.54 v.
• Longer shelf life than acidic dry cells and rechargeable; little corrosion of zinc.
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Lead Storage Battery• Six cells in series.
• Electrolyte = 6 M H2SO4.
• Anode = Pb.Pb(s) + SO4
2-(aq) PbSO4(s) + 2 e-
• Cathode = Pb coated with PbO2.
• PbO2 is reduced.PbO2(s) + 4 H+(aq) + SO4
2-(aq) + 2 e- PbSO4(s) + 2 H2O(l)
• Cell voltage = 2.09 v.
• Rechargeable, heavy.
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Fuel Cells• Like batteries in which reactants are constantly being
added.So it never runs down!
• Anode and cathode both Pt-coated metal.• Electrolyte is OH– solution.• Anode reaction: 2 H2 + 4 OH– → 4 H2O(l) + 4 e-.• Cathode reaction: O2 + 4 H2O + 4 e- → 4 OH–.
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Nonspontaneous Redox Reaction
• The reverse of a spontaneous reaction is nonspontaneous.
• To get it to run, an outside energy source must be supplied.
• Nonspontaneous redox reactions can be made to work by using a battery to force the electrons to flow in the nonspontaneous direction.
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Electrolysis • Electrolysis is the process of using
electricity to break a compound apart.
• Electrolysis is done in an electrolytic cell.
• Electrolytic cells can be used to separate elements from their compounds.Generate H2 from water for fuel cells.Recover metals from their ores.
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Electrolytic Cell• The + terminal of the battery = anode.• The - terminal of the battery = cathode.• Cations attracted to the cathode; anions attracted to the
anode.• Cations pick up electrons from the cathode and are
reduced; anions release electrons to the anode and are oxidized.
• In electroplating, the work piece is the cathode.Cations are reduced at the cathode and plate onto the surface.The anode is made of the plate metal, the anode oxidizes and
replaces the metal cations lost from the solution.
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Electrolytic Cell—Electroplating
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Corrosion• Corrosion is the spontaneous oxidation of a
metal by chemicals in the environment.• Since many materials we use are active metals,
corrosion can be a very big problem.
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Preventing Corrosion• One way to reduce or slow corrosion is to coat
the metal surface to keep it from contacting corrosive chemicals in the environment.Paint.Some metals, like Al, form an oxide that strongly
attaches to the metal surface, preventing the rest from corroding.
• Another method to protect one metal is to attach it to a more reactive metal that is cheap.Sacrificial electrode.