Transcript Chapter 22 REDOX

Chapter 22 REDOX
Chapter 22 REDOX
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The Meaning of Oxidation and Reduction
Oxidation Numbers
Balancing Redox Equations
Ch 22.1 The Meaning
of Oxidation and Reduction
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Oxygen in Redox Reactions
Electron Transfer in Redox Reactions
Corrosion
Oxygen in Redox Reactions
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Oxidation – the combination of an element
with oxygen to produce oxides
Oxygen in Redox Reactions
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Burning
Oxygen in Redox Reactions
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Bleaching
Oxygen in Redox Reactions
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Rusting
Reduction
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Reduction – the loss
of oxygen from a
compound
Redox Reactions
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Reduction and Oxidation always occur
together
2Fe2O3(s) + 3C(s)  4Fe(s) + 3CO2(g)
reduction oxidation
Electron Transfer in Redox
Reactions
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Oxidation – loss of electrons, gain oxygen
Reduction – gain of electrons, loss of
oxygen
“LEO the lion goes GER”
LEO – Lose electrons oxidation
GER – Gain electrons reduction
Electron Transfer in Redox
Reactions
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Oxidation: Mg  Mg2+ + 2e
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Loss of electrons
Reduction: S + 2e-  S2
Gain of electrons
Corrosion
Corrosion
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2Fe(s) + O2(g) + 2H2O(l)  2Fe(OH)2(s)
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4Fe(OH)2(s) + O2(g) + 2H2O(l)  4Fe(OH)3(s)
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Corrosion of iron
Corrosion
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Some metals completely corrode
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Some metals form a protective coating
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Iron
Aluminum
Some metals do not corrode at all
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Gold
Chapter 22.3
Balancing Redox Reactions
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Identifying Redox Reactions
Using Oxidation Number Changes
Using Half Reactions
Identifying Redox Reactions
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Two types of reactions:
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REDOX – electrons are transferred
Everything else: single replacement, double
replacement, combustion, ….
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NO transfer of electrons
Identifying Redox Reactions
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REDOX – the oxidation number of an
element changes
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N2(g) + O2(g)  2NO(g)
Using Oxidation Number
Changes
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Fe2O3(s) + CO(g)  Fe(s) + CO2(g)
Step 1 – Assign oxidation numbers to all
atoms in the equation
Fe2O3(s) + CO(g)  Fe(s) + CO2(g)
+3 -2 +2 –2
0
+4 -2
Using Oxidation Number
Changes
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Step 2 – Identify which atoms are oxidized
and reduced
Fe2O3(s) + CO(g)  Fe(s) + CO2(g)
+3 -2 +2 -2
0
+4 -2
Iron – reduced, Carbon - oxidized
Using Oxidation Number
Changes
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Step 3 – Use a bracket line to connect the
atoms undergoing oxidation and one to
connect the lines undergoing reduction
+2
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Fe2O3(s) + CO(g)  Fe(s) + CO2(g)
-3
Using Oxidation Number
Changes
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Make the total increase in oxidation
number equal to the total decrease in
oxidation number by using appropriate
coefficients
3 x (+2) = 6
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Fe2O3(s) + CO(g)  Fe(s) + CO2(g)
2 x (-3) = - 6
Using Oxidation Number
Changes
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Step 5 – Finally make sure the equation is
balanced for both atoms and charge
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Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g)
FRH – Flameless Ration Heater
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Mg + H2O  Mg(OH)2 + H2 + Heat
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Problem: Mg forms a coating from corrosion –
MgO – which is not water soluble, prevents the
above reaction from happening
Solution: Add NaCl and Fe to the mix, breaks
down the MgO and allows the reaction to
happen
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Chapter 23 Electrochemistry
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Electrochemical Cells
Half Cells and Cell Potentials
Electrolytic Cells
Electrochemical Cells
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The Nature of Electrochemical Cells
Voltaic Cells
Dry Cells
Lead Storage Batteries
Fuel Cells
The Nature of Electrochemical
Cells
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Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
The Nature of Electrochemical
Cells
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The zinc bar becomes copper plated
Zinc loses electrons and dissolves slowly
Copper gains electrons and becomes a
solid
Oxidation: Zn(s) Zn2+(aq) + 2eReduction: Cu2+(aq) + 2e- Cu(s)
The Nature of Electrochemical
Cells
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Reference Table J
Look at any two metals, the metal that is
higher on the table is the one that is more
readily oxidized
Electrochemical Cell
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Any device that converts chemical energy
into electrical energy or electrical energy
into chemical energy
REDOX reactions must occur
If an electrochemical cell is to be used for
electrical energy, the two half reactions
must physically be separated
Voltaic Cell
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Alessandro Volta (1745 – 1827)
First electrochemical cell
Voltaic Cell
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Convert chemical energy into electrical
energy
Half Cell – part of a voltaic cell, consists of
a metal rod in a solution of ions
Salt Bridge – Separates half cells, tube
containing a strong electrolyte (can also
use a porous plate)
Voltaic Cell
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Anode – the anode where oxidation
occurs
Cathode – the cathode where reduction
occurs
Dry Cell
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A voltaic cell in which the electrolyte is a
paste
Alkaline Battery
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Improved dry cell, the Zinc electrode
doesn’t corrode as fast
Lead Storage Batteries
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A group of cells connected together
Fuel Cell
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Voltaic cell in which a fuel substance
undergoes oxidation
Do not have to be recharged
Oxidation: 2H2(g) + 4OH-(aq)  4H2O(l) + 4eReduction: O2(g) + 2H2O(l) + 4e-  4OH-(aq)
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http://automobiles.honda.com/fcx-clarity/
http://www.pbs.org/wgbh/nova/sciencenow
/3210/01.html
Hydrogen Refueling Stations
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The diagram shows a voltaic cell with copper and
aluminum electrodes immediately after the external
circuit is completed.
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1 Balance the redox equation using the smallest wholenumber coefficients. [1]
2 As this voltaic cell operates, the mass of the Al(s)
electrode decreases. Explain, in terms of particles, why
this decrease in mass occurs. [1]
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Answers
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3 Cu2+ (aq) + 2 Al(s)  3 Cu(s) + 2 Al3+ (aq)
Aluminum particles are losing electrons
and becoming aluminum ions that are
entering the solution.
It allows migration of ions, maintains
neutrality, prevents polarization
Electrolytic Cells
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An electrochemical cell used to cause
chemical change through the application
of electrical energy (electrical energy is
added)
Differences
Voltaic (Galvanic) Cells
Electrolytic Cells
Flow of electrons is
spontaneous
Flow of electrons is
pushed by an outside
power source
Anode negative
Cathode positive
Anode positive
Cathode negative
Similarities
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Voltaic (Galvanic) Cells and Electrolytic
Cells
Electrons flow from anode to cathode
Reduction – cathode
Oxidation – anode
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Electroplating is an electrolytic process
used to coat metal objects with a more
expensive and less reactive metal. The
diagram below shows an electroplating
cell that includes a battery connected to a
silver bar and a metal spoon. The bar and
spoon are submerged in AgNO3(aq).
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Explain why AgNO3 is a better choice than AgCl
for use in this electrolytic process. [1]
Explain the purpose of the battery in this cell. [1]
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Acceptable responses include, but are not
limited to:
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Silver nitrate produces more ions than silver chloride
in water.
AgNO3 readily dissolves in H2O; AgCl dissolves only
slightly in H2O.
Acceptable responses include, but are not
limited to:
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The battery provides the electrical energy necessary
for the reaction to occur.
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The apparatus shown in the diagram
consists of two inert platinum electrodes
immersed in water. A small amount of an
electrolyte, H2SO4, must be added to the
water for the reaction to take place. The
electrodes are connected to a source that
supplies electricity.
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What type of electrochemical cell is shown? [1]
What particles are provided by the electrolyte
that allow an electric current to flow? [1]
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Electrolytic or electrolysis.
Acceptable responses include, but are not
limited to:
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Ions, charged particles, H3O+, SO42–
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Because tap water is slightly acidic, water
pipes made of iron corrode over time, as
shown by the balanced ionic equation
below:
2Fe + 6H+  2Fe3+ + 3H2
Explain, in terms of chemical reactivity,
why copper pipes are less likely to corrode
than iron pipes. [1]
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Acceptable responses include, but are not
limited to:
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Copper is less reactive than iron.
Cu below H2 on Table J