Chapter 22 REDOX. The Meaning of Oxidation and Reduction Oxidation Numbers Balancing Redox...
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Transcript of Chapter 22 REDOX. The Meaning of Oxidation and Reduction Oxidation Numbers Balancing Redox...
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Chapter 22 REDOX
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Chapter 22 REDOX
The Meaning of Oxidation and Reduction Oxidation Numbers Balancing Redox Equations
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Ch 22.1 The Meaning of Oxidation and Reduction
Oxygen in Redox Reactions Electron Transfer in Redox Reactions Corrosion
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Oxygen in Redox Reactions
Oxidation – the combination of an element with oxygen to produce oxides
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Oxygen in Redox Reactions
Burning
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Oxygen in Redox Reactions
Bleaching
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Oxygen in Redox Reactions
Rusting
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Reduction
Reduction – the loss of oxygen from a compound
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Redox Reactions
Reduction and Oxidation always occur together
2Fe2O3(s) + 3C(s) 4Fe(s) + 3CO2(g)
reduction oxidation
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Electron Transfer in Redox Reactions
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
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Electron Transfer in Redox Reactions
Oxidation: Mg Mg2+ + 2e-
Loss of electrons Reduction: S + 2e- S2-
Gain of electrons
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Corrosion
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Corrosion
2Fe(s) + O2(g) + 2H2O(l) 2Fe(OH)2(s)
4Fe(OH)2(s) + O2(g) + 2H2O(l) 4Fe(OH)3(s)
Corrosion of iron
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Corrosion
Some metals completely corrode Iron
Some metals form a protective coating Aluminum
Some metals do not corrode at all Gold
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Chapter 22.3 Balancing Redox Reactions
Identifying Redox Reactions Using Oxidation Number Changes Using Half Reactions
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Identifying Redox Reactions
Two types of reactions: REDOX – electrons are transferred Everything else: single replacement, double
replacement, combustion, …. NO transfer of electrons
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Identifying Redox Reactions
REDOX – the oxidation number of an element changes N2(g) + O2(g) 2NO(g)
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Using Oxidation Number Changes
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
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Using Oxidation Number Changes
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
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Using Oxidation Number Changes
Step 3 – Use a bracket line to connect the atoms undergoing oxidation and one to connect the lines undergoing reduction
Fe2O3(s) + CO(g) Fe(s) + CO2(g)
-3
+2
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Using Oxidation Number Changes
Make the total increase in oxidation number equal to the total decrease in oxidation number by using appropriate coefficients
Fe2O3(s) + CO(g) Fe(s) + CO2(g)
3 x (+2) = 6
2 x (-3) = - 6
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Using Oxidation Number Changes
Step 5 – Finally make sure the equation is balanced for both atoms and charge
Fe2O3(s) + 3CO(g) 2Fe(s) + 3CO2(g)
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FRH – Flameless Ration Heater
Mg + H2O Mg(OH)2 + H2 + Heat
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
Electrochemical Cells Half Cells and Cell Potentials Electrolytic Cells
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Electrochemical Cells
The Nature of Electrochemical Cells Voltaic Cells Dry Cells Lead Storage Batteries Fuel Cells
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The Nature of Electrochemical Cells
Zn(s) + Cu2+(aq) Zn2+
(aq) + Cu(s)
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The Nature of Electrochemical Cells
The zinc bar becomes copper plated Zinc loses electrons and dissolves slowly Copper gains electrons and becomes a
solid Oxidation: Zn(s) Zn2+
(aq) + 2e-
Reduction: Cu2+(aq) + 2e- Cu(s)
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The Nature of Electrochemical Cells
Reference Table J Look at any two metals, the metal that is
higher on the table is the one that is more readily oxidized
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Electrochemical Cell
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
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Voltaic Cell
Alessandro Volta (1745 – 1827) First electrochemical cell
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Voltaic Cell
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)
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Voltaic Cell
Anode – the anode where oxidation occurs
Cathode – the cathode where reduction occurs
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Dry Cell
A voltaic cell in which the electrolyte is a paste
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Alkaline Battery
Improved dry cell, the Zinc electrode doesn’t corrode as fast
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Lead Storage Batteries
A group of cells connected together
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Fuel Cell
Voltaic cell in which a fuel substance undergoes oxidation
Do not have to be recharged Oxidation: 2H2(g) + 4OH-
(aq) 4H2O(l) + 4e-
Reduction: 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
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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.
1 Balance the redox equation using the smallest whole-number 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]
3 Explain the function of the salt bridge. [1]
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Answers
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
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Electrolytic Cells
An electrochemical cell used to cause chemical change through the application of electrical energy (electrical energy is added)
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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
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Similarities
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: 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: 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: Ions, charged particles, H3O+, SO4
2–
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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: Copper is less reactive than iron. Cu below H2 on Table J