Chapter 16 Oxidation and Reduction

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Chapter 16OxidationandReduction
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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 ReductionOne 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 OxidationReduction

• 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 Na2O (s)
• 2 Na(s) Cl2(g) ? 2 NaCl(s)

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Oxidation and ReductionAnother 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 NaCl(s) Na ? Na 1 e
oxidation Cl2 2 e ? 2 Cl reduction
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OxidationReduction

• 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 NaCl(s) Na is oxidized, Cl
is reduced Na is the reducing agent, Cl2 is the
oxidizing agent
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Practice Identify the Element being Oxidized
and Element being Reduced and the Oxidizing and
Reducing Agents

• 2 Mg(s) O2(g) ? 2 MgO(s)
• Fe(s) Cl2(g) ? FeCl2(s)
• Zn(s) Fe2(aq) ? Zn2(aq) Fe(s)

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Practice Identify the Element being Oxidized
and Element being Reduced and the Oxidizing and
Reducing Agents

• 2 Mg(s) O2(g) ? 2 MgO(s)
• Fe(s) Cl2(g) ? FeCl2(s)
• Zn(s) Fe2(aq) ? Zn2(aq) Fe(s)

Mg is oxidized, O is reduced Mg is the reducing
agent, O2 is the oxidizing agent
Fe is oxidized, Cl is reduced Fe is the reducing
agent, Cl2 is the oxidizing agent
Zn is oxidized, Fe is reduced Zn is the reducing
agent, Fe2 is the oxidizing agent
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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
• even though 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
• free elements have an oxidation state 0
• Na 0 and Cl2 0 in 2 Na(s) Cl2(g)
• monatomic ions have an oxidation state equal to
  their charge
• Na 1 and Cl -1 in NaCl
• (a) the sum of the oxidation states of all the
  atoms in a compound is 0
• Na 1 and Cl -1 in NaCl, (1) (-1) 0

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Rules for Assigning Oxidation States

• (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
• (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

• in their compounds, nonmetals have oxidation
  states according to the table below
• nonmetals higher on the table take priority

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Practice Assign an Oxidation State to each
Element in the following

• Br2
• K
• LiF
• CO2
• SO42-
• Na2O2

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Practice Assign an Oxidation State to each
Element in the following

• Br2 Br 0, (Rule 1)
• K K 1, (Rule 2)
• LiF Li 1, (Rule 4a) F -1, (Rule 5)
• CO2 O -2, (Rule 5) C 4, (Rule 3a)
• SO42- O -2, (Rule 5) S 6, (Rule 3b)
• Na2O2 Na 1, (Rule 4a) O -1, (Rule 3a)

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Oxidation and ReductionA Better Definition

• oxidation occurs when an atoms oxidation state
  increases during a reaction
• reduction occurs when an atoms oxidation state
  decreases during a reaction

CH4 2 O2 ? CO2 2 H2O -4 1
0 4 2 1 -2
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Identify the Oxidizing and Reducing Agents in
Each of the Following

• 3 H2S 2 NO3 2 H 3 S 2 NO 4 H2O
• MnO2 4 HBr MnBr2 Br2 2 H2O

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Identify the Oxidizing and Reducing Agents in
Each of the Following
ox ag
red ag

• 3 H2S 2 NO3 2 H 3 S 2 NO 4 H2O
• MnO2 4 HBr MnBr2 Br2 2 H2O

1 -2 5 -2 1 0
2 -2 1 -2
red ag
ox ag
4 -2 1 -1 2 -1 0
1 -2
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Balancing Redox Reactions

• assign oxidation states and determine element
  oxidized and element reduced
• separate into oxidation reduction
  half-reactions
• balance half-reactions by mass
• first balance atoms other than O and H
• then balance O by adding H2O to side that lacks O
• finally balance H by adding H to side that lacks
  H

Fe2 ? Fe3
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Balancing Redox Reactions

• balance each half-reaction with respect to charge
  by adjusting the numbers of electrons
• electrons on product side for oxid.
• electrons on reactant side for red.
• balance electrons between half-reactions
• add half-reactions, canceling electrons and
  common species
• Check

Fe2 ? Fe3 1 e-
MnO4 8H 5 e- ? Mn2 4H2O
x 5
5 Fe2 MnO4 8H ? Mn2 4H2O 5 Fe3
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Practice Balance the Following EquationCu
I2 ? Cu2 I
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Practice Balance the Following EquationCu
I2 ? Cu2 I
1
0
2
-1
oxid
red
ox Cu ? Cu2
red I2 ? I
ox Cu ? Cu2
red I2 ? 2 I
ox Cu ? Cu2 1 e-
red I2 2 e- ? 2 I
ox Cu ? Cu2 1 e- x 2
red I2 2 e- ? 2 I
2 Cu I2 ? 2 Cu2 I2
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Practice Balance the Following Equation I
Cr2O72- ? Cr3 I2
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Practice Balance the Following Equation I
Cr2O72- ? Cr3 I2
6
0
3
-1
-2
oxid
red
ox I ? I2
red Cr2O72 ? Cr3
ox 2 I ? I2
red Cr2O72 ? 2 Cr3
red Cr2O72 ? 2 Cr3 7 H2O
red Cr2O72 14H? 2Cr3 7H2O
ox 2 I ? I2 2e-
red Cr2O72 14H 6e- ? 2Cr3 7H2O
ox 2 I ? I2 2e-x3
red Cr2O72 14H 6e- ? 2Cr3 7H2O
Cr2O72 14 H 6 I ? 2 Cr3 7 H2O 3 I2
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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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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

Zn Fe2 Fe Zn2
Cu Fe2 no reaction
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Mg is above Cu on the Activity Series
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Predict the Products Balance the Equation

• Mg H3PO4
• Cu H2SO4
• Al Fe2

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Predict the Products Balance the Equation

• 3 Mg 2 H3PO4 Mg3(PO4)2 3 H2
• Cu H2SO4 no reaction
• 2 Al 3 Fe2 2 Al3 3 Fe

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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

• 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
• Cathode
• electrode 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 number of electrons that pass a point each
second is called the current of the electricity.
The amount of water that passes a point each
second is called the current of the river.
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Voltage
Voltage is the force pushing the electrons down
the wire.
Gravity is the force pulling the water down the
river.
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Dead Battery
As the reaction proceeds, the reactants
get consumed and the voltaic cell dies. The
current decreases until electrons can no
longer flow through the wire.
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LeClanche 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

• 6 cells in series
• electrolyte 6 M H2SO4
• anode Pb
• Pb(s) SO42-(aq) PbSO4(s) 2 e-
• cathode Pb coated with PbO2
• PbO2 is reduced
• PbO2(s) 4 H(aq) SO42-(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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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

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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 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 workpiece is the cathode
• cations are reduced at cation and plate to the
  surface
• the anode is made of the plate metal, the anode
  oxidizes and replaces the metal cations in the
  solution

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Electrolytic Cell - Electroplating