Unit 1: Electrochemistry

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Unit 1 Electrochemistry

• Redox Reactions

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Redox Reactions

• What do the combustion of a hydrocarbon, the
  smelting of metallic ores, the rusting of an
  automobile, the electrolysis of water, the silver
  plating of cutlery and the functioning of a
  battery all have in common?  They are all
  examples of a unique type of chemical reaction
  called an oxidation-reduction reaction or redox
  for short. 
• The distinguishing characteristic of redox
  reactions is the apparent transfer of one or more
  electrons from one reactant to another reactant
  as the system proceeds from reactant to product. 
  This unit is about the comings and goings of
  electrons in chemical reactions.

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Oxidation

• Oxidation was a term used to describe any
  reaction in which oxygen (O2) was involved as a
  reactant.  Combustion reactions such as
• CH4(g) O2(g) ----- CO2(g) 2 H2O(g)
• and corrosion reactions such as
• 4 Fe(s) 3 O2(g) ----- 2 Fe2O3(s)
• were called oxidation reactions. 
• It later became obvious that similar reactions
  could take place whether oxygen was present or
  not but the term oxidation was retained.

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Reduction

• Reduction was used to refer to the smelting
  process, whereby tons of metallic ore were
  reduced to a much smaller quantity of pure metal
  during the extraction.  The production of pure
  tin from tin ore is an example
• 2 SnO(s) C(s) ----- 2 Sn(s) CO2(g)
• As knowledge about these reactions grew, it
  became obvious that there was a more fundamental
  similarity between these types of reactions than
  the simple reduction of the mass of an ore to the
  smaller mass of a refined metal but again the
  term reduction was retained.

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• These terms are normally used to describe the
  processes of electron loss and gain
• Reduction - a process whereby a substance gains
  electrons in a redox system.Oxidation - a
  process whereby a substance loses electrons in a
  redox system.
• To avoid confusion, memorize the following
  mnemonic device
•     LEO the lion says GER
• where  LEO Loss of electrons is oxidation
              GER Gain of electrons is reduction
  

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Redox Reactions

• A Redox reaction is one in which both processes
  take place simultaneously, the electrons lost by
  one reactant are the same electrons gained by the
  other.
• Having said all this, we often refer to these
  processes as though they could operate in
  isolation.  The reaction that symbolically
  describes this is called a half-reaction. 
• Here are some examples

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• Here are some examples of oxidation half
  reactions in which substances lose one or more
  electrons.  Note that the electrons are written
  on the product side of the equation
• a)  Al(s) ----- Al3(aq) 3e- b)  2 I-(aq)
  ----- I2(s) 2e- c)  2 H2O(l) ----- 4 H(aq)
  O2(g) 4e- d)  Fe2(aq) ----- Fe3(aq) e-

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• Likewise it is sometimes desirable to focus on
  the reduction process to the exclusion of
  oxidation.  Here are some examples of reduction
  half reactions in which familiar substances gain
  one or more electrons.  Note that the electrons
  are written on the reactant side of the equation
  
• a)  Cu2(aq) 2e- ----- Cu(s) b)  Ag(aq)
  e- ----- Ag(s) c)  Fe3(aq) e- ----- Fe2(aq)
  d)  2 H2O(l) 2e- ----- H2(g) 2 OH-(aq)

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Oxidizing Agents and Reducing Agents

• Agent is defined as something that performs
  actions or produces an affect. 
• Any substance which takes electrons from another
  substance is causing the oxidation of that other
  substance.  There are three equivalent ways of
  describing an oxidizing agent (OA)
• And, any substance which gives electrons to
  another substance is causing the reduction of
  that other substance.  There are three equivalent
  ways of describing a reducing agent (RA)

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Oxidizing Agent

• a substance which causes the oxidation of another
  substance
• a substance that itself undergoes reduction
• a substance whose action is represented by a
  reduction half reaction
•         ex. Cu2(aq) 2 e- ----- Cu(s)
• The Cu2 (aq) gained two electrons from another
  substance.  A half-reaction does not tell us what
  that other substance is or what happened to it
  but we assume that it is out there somewhere,
  short two electrons.  That other substance has
  been oxidized by the copper(II) ion.  The
  copper(II) ion is therefore an oxidizing agent. 
  Note that in the process of oxidizing another
  substance, it has itself become reduced, so its
  action is expressed in a reduction half-reaction.
  
•  

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Reducing Agent 

• a substance which causes the reduction of another
  substance
• a substance that itself undergoes oxidation
• a substance whose action is represented by an
  oxidation half reaction
•         ex.  Al(s) ----- Al3(aq) 3 e-
• The neutral aluminum atom (Al) lost three
  electrons to another substance.  Once again the
  half- reaction does not tell us what or where
  that other substance is but we assume it is out
  there somewhere trying to cope with three extra
  electrons.  It has been reduced by the aluminum
  atom.  The aluminum atom is therefore a reducing
  agent.  Note that in the process of reducing the
  other substance, it has itself become oxidized,
  so its action is expressed in an oxidation
  half-reaction.

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Half reactions can be read in both directions

• As we read from left to right in the following
  reaction
• Al3(aq) 3 e- Al(s)
• Al3(aq) is an oxidizing agent undergoing
  reduction to become Al(s).  However, Al(s) is a
  reducing agent undergoing oxidation to become
  Al3(aq) as we read the reverse reaction from
  right to left.  In other words, a reducing agent
  becomes an oxidizing agent and vice versa.  Like
  oxidation and reduction, the terms oxidizing
  agent and reducing agent can be easily confused
  with each other.
• To avoid confusion, remember the word OAR (as in
  boat oar)
• Read left to right it can be used to remember
  that         Oxidizing Agents undergo Reduction
  
• Read right to left it can be used to remember
  that       Reducing Agents undergo Oxidation

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The Reduction Potential Table
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Combining Oxidation and Reduction

• Using the Reduction Potential Table
• Since neither oxidation nor reduction can occur
  in isolation, the Oxidizing and Reducing Agents
  Table on page 9 of your Data Booklet is a useful
  tool for learning how to glue the two halves of a
  Redox reaction together.  The table can be used
  to
• 1.  predict the most likely redox half-reactions
  in a given chemical system, 2.  write balanced
  net ionic equations for the reactions, and 3. 
  determine whether or not the reaction will be
  spontaneous.

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Example 1

• Predict the most likely redox reaction to occur
  when a strip of solid chromium is placed in an
  aqueous solution of gold (III) nitrate, balance
  the equation for the reaction and determine
  whether or not the reaction will be spontaneous.
• Solution
• There are 5 steps needed to write this redox
  equation in an orderly fashion...

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Step 1

• List all the chemical species present in the
  beaker.  Use the Table to identify all the
  possible oxidizing agents (OA) and all the
  possible reducing agents (RA).  Remember that
  oxidizing agents are listed on the left side of
  the arrow in the table and reducing agents are
  listed on the right side.  Note that some
  chemical species can act as an oxidizing agent
  and as a reducing agent.
• (OA)              (OA)             
  (OA) Cr(s)         Au3(aq)        
  NO3-(aq)         H2O(l) (RA)                     
                                 (RA)

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Step 2

• Use the Table to identify which of the oxidizing
  agents is strongest (SOA) and write its half
  reaction.  Copy it right out of the table. 
  Remember that the strongest oxidizing agent is
  listed on the left side nearest the top of the
  table.
• So the SOA is Au3(aq) and its half reaction is
  
• Au3(aq) 3 e- ------ Au(s)
• Important  You might think the NO3-(aq) is the
  best choice for SOA but you would be wrong!!! 
  Complex or polyatomic ions can only serve
  effectively as oxidizing agents in acidic or
  basic solutions.  There is no acid or base
  mentioned here.  Neither an acid formula nor the
  word "acidified" appears in the question. 
  NO3-(aq) in the absence of hydrogen ions (acid)
  cannot compete, and will serve only as a
  spectator ion.

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Step 3

• Use the Table to identify which of the potential
  reducing agents is the strongest (SRA) and write
  its half reaction. Remember that the strongest
  reducing agent is listed on the right side of the
  arrows nearest the bottom of the table.
• The SRA is Cr(s) and its half reaction is
• Cr(s) ------ Cr2(aq) 2 e-

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Step 4

• Write the balanced net ionic equation for the
  reaction.  Multiply the half reactions by the
  smallest whole numbers to make the gain of
  electrons equal the loss of electrons, then add
  the half reactions together
•           
• 2( Au3(aq) 3 e- ------ Au(s) )
              3( Cr(s) ------ Cr2(aq) 2 e- )
  2 Au3(aq) 3 Cr(s) ------ 2 Au(s) 3 Cr2(aq)

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Step 5

• Determine whether or not the reaction is
  spontaneous.  This is most simply done be
  comparing the relative position of the OA and RA
  on the table.  If the oxidizing agent is listed
  higher (nearer the top of the page) in the Table
  than the reducing agent, then the reaction is
  spontaneous.  If the oxidizing agent is listed
  lower (nearer the bottom of the page) in the
  Table than the reducing agent, then the reaction
  is nonspontaneous.  If the oxidizing and reducing
  agents are at the same level the reaction is also
  nonspontaneous.
• Since the Au3(aq) (SOA) is listed higher in the
  Table than is the Cr(s) (SRA) the reaction is
  spontaneous.  Put your fingers on them in the
  table and see for yourself... your left pointing
  finger should be higher that your right pointing
  finger.  (downhill slope)

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Example 2

• Predict the most likely redox reaction to occur
  when sulfuric acid and aqueous solutions of
  sodium permanganate and tin (II) chloride are
  mixed together in the same beaker.  Balance the
  equation for the reaction and determine whether
  or not it will be spontaneous.  This question has
  just about every "wrinkle" in it I could find.

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Step 1

• OA             OA?        OA       
  spectator         OA                          OA
  Na(aq)     MnO4-(aq)     H(aq)    
  HSO4-(aq)     Sn2(aq)     Cl-(aq)     H2O(l)
                                                   
                          SRA          
  RA           RA
• Before you proceed to complete steps 2 to 5 on
  your own, note both MnO4-(aq) and H(aq) are
  potential oxidizing agents, but the combination
  of the two (acidified permanganate) produces the
  strongest oxidizing agent. 

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In case you missed it

• 1. Sulfuric acid is a strong acid and should be
  shown in dissociated form.
• 2. Assume that any species not listed on the
  table is a spectator species.
• 3. The MnO4-(aq), like most polyatomic ions,
  must be in acidic solution (H(aq)) in order to
  act as a strong oxidizing agent.  Otherwise, they
  are merely spectators.

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Practice

• Now work on the questions in the worksheet
  package on page M11, M12 and Worksheet 54.
• After this is complete, you should now be able to
  complete the Chapter 12 Assignment 1 that will
  be due on ___________.