Reduction – Oxidation Reactions “REDOX”
• Examples of Redox Reactions:
 Formation, decomposition, combustion, single replacement, cellular
respiration, photosynthesis,
(NOT double
replacement)
Predicting Redox Reactions 13.2
 REMEMBER a redox reaction is a transfer of valence
electrons from one substance to another.
 Chemists use the analogy of a tug-of-war to describe
the competition for electrons
 In a reaction, each entity pulls on the same electrons
 If one entity is able to pull electrons away from the other,
a spontaneous reaction occurs
 If one entity is unable to pull electrons away from the
other, no reaction occurs (non-spontaneous)
How do we know…?
 Without mixing all possible reactants and observing any
evidence of reaction, how can we predict if a reaction will
occur? (spontaneous or nonspontaneous)??
 If a reaction occurs, what will be produced??
 The answers to these questions cannot be obtained easily
using redox theory.
 By observing many successful and unsuccessful reactions,
patterns emerge and empirical generalizations can be made.
 But first before looking at patterns we need to understand
some terms: OA (oxidizing agent) and RA (reducing agent)…
Oxidizing and Reducing
Agents
 In any redox reaction, an electron transfer occurs:
 One reactant is oxidized, and one reactant is reduced.
Terms to identify e- transfer.
 Rather than saying “the reactant that is oxidized” and
“the reactant that is reduced,” chemists use the terms
reducing agent (RA) and oxidizing agent (OA)
Reducing Agent (RA)
Oxidizing Agent (OA)
causes reduction
causes oxidization
is oxidized
is reduces
loses electron(s)
gains electron(s)
(usually more stable metals) (usually reactive nonmetals)
Note: Remember when an atom becomes more + it is losing electrons.
 A reducing agent causes reduction by donating
(losing) electrons. During this process it is oxidized.
 An oxidizing agent causes oxidation by removing
(gaining) electrons. During this process it is reduced.
 It is important to note that oxidation and reduction are
processes, and oxidizing agents and reducing
agents are substances.
Note: From ancient times “reduced” meant to take a substance (compound) and
“reduce” it to its pure form. Reduced = is gaining electrons.
Also anciently the term “oxidizing agent” was thought to only oxygen but then
found that others (eg. Halogens) can also oxidize/corrode metals.
Redox Terms

Silver ions were reduced to silver metal by reaction with copper metal. Simultaneously, copper
metal was oxidized to copper(II) ions by reaction with silver ions.

If Ag+(aq) is reduced it is the:

If Cu(s) is oxidized it is the:
It is important to note that oxidation and reduction are processes,
and oxidizing agents and reducing agents are substances.
REDOX Reactions … so far
Reduction
•
Historically, the formation of a metal from its
“ore” (or oxide)
▫ I.e. nickel(II) oxide is reduced by
hydrogen gas to nickel metal
Oxidation
•
▫
NiO(s) + H2(g)  Ni(s) + H2O(l)
Ni
+2

Historically, reactions with oxygen
I.e. iron reacts with oxygen to produce
iron(III) oxide
4 Fe(s) + O2(g)  Fe2O3(s)
Nio
Fe 0

Fe+3
•
A gain of electrons occurs (so the entity
becomes more negative)
•
A loss of electrons occurs (so the entity
becomes more positive)
•
Electrons are shown as the reactant in the
half-reaction
•
Electrons are shown as the product in the halfreaction
•
A species undergoing reduction will be
responsible for the oxidation of another entity
– and is therefore classified as an oxidizing
agent (OA)
•
A species undergoing oxidation will be
responsible for the reduction of another entity –
and is therefore classified as an reducing agent
(RA)
Redox Terms
 Summary so far:
 The substance that is reduced (gains electrons) is
also known as the oxidizing agent
 The substance that is oxidized (loses electrons) is
also knows as the reducing agent
(Is Reduced)
(Is Oxidized)
• Question: If a substance is a very strong oxidizing agent, what
does this mean in terms of electrons?
The substance has a very strong attraction for electrons.
• Question: If a substance is a very strong reducing agent, what
does this mean in terms of electrons?
The substance has a weak attraction for its electrons, which are easily removed
Development of a Redox
Table
 In the past we have generally assumed that all single
replacement reactions are spontaneous. But now we
can’t assume… so…
 How do you know when a chemical reaction will occur
spontaneously without actually doing the reaction?
 Experiments and a redox table.
Redox Table
 A reaction is considered spontaneous if it occurs on its own
 A reduction ½ reaction table is useful in predicting the spontaneity of
a reaction
 Reduction Tables show reduction ½ reactions in the forward direction,
therefore all the reactants will be oxidizing agents
 If we list the OA’s from (P.569 look at table 1) in decreasing order of
strength, we create a reduction ½ reaction table:
WRA
SOA
WOA
Ag+(aq) + 1 e-  Ag(s)
Cu2+(aq) + 2 e-  Cu(s)
Pb2+(aq) + 2 e-  Pb(s)
Zn2+(aq) + 2 e-  Zn(s)
SRA
 Strongest oxidizing agent at the top left
 Strongest reducing agent at the bottom right
 For metal ions (oxidizing agents), the half-reactions are
read from left to right in the table.
 For metal atoms (reducing agents), the half-reactions are
read from right to left.
Reading a Redox Table
 The table lists the relative strengths of oxidizing and
reducing agents.
 Listed as reductions (from left to right) in the form:
 The strongest oxidizing agent is listed at the top left
 The strongest reducing agent is listed at the bottom
right
Summary
 An oxidizing agent causes oxidation by removing (gaining)
electrons from substance in a redox reaction. In the process
the oxidizing agent is reduced.
 A reducing reagent promotes reduction by donating (loosing)
electrons to another substance in a redox reaction. In the
process the reducing agent is oxidized.
 A table of relative strengths of oxidizing and reducing agents –
known as a redox table – is, by convention, listed as reductions
(from left to right) in the form
, with the strongest
oxidizing agent at the top left and strongest reducing agent at
the bottom right of the table.
Homework…
 p.571 #1-3,6,7
 The redox spontaneity rule states that a spontaneous
redox reaction occurs only if the oxidizing agent is above
the reducing agent in the redox table.
Thankfully for us…
 Chemists have analyzed, experimented, and collected
enough data to be able to produce an extended redox
table of oxidizing and reducing agents…. And that table
is found in the back of your text book (Appendix 1 p.
828 or in Data booklets).
 You can use this table as a reference when trying to
predict spontaneous redox reactions.
Appendix 1 p. 828 or in Data
booklets
This picture is from your
data booklet reduction
½ reaction table
 The relative position of a pair of oxidizing agents and
reducing agents indicates whether a reaction will be
spontaneous.
Building Redox Tables #2

Example 2: Use the following information to create a table of reduction ½ reactions
OA
3 Co 2+ (aq) + 2 In(s)  2 In 3+ (aq) + 3 Co(s)
OA
Cu 2+ (aq) + Co(s)  Co 2+ (aq) + Cu(s)
OA
Cu 2+ (aq) + Pd(s)  no reaction
SOA
Pd2+(aq)
Cu2+(aq)
Co2+(aq)
In3+(aq)
+ 2 e-  Pd(s)
+ 2 e-  Cu(s)
+ 2 e-  Co(s)
+ 3 e-  In(s)
SRA
Building Redox Tables #3

Example 3: Use the following information to create a table of reduction ½ reactions
OA
2 A 3+ (aq) + 3 D(s)  3 D2+ (aq) + 2 A(s)
A3+
OA
G
+
(aq)
+
D2+(aq)
D(s)  no reaction
D(s)
E(s)
OA
3 D 2+ (aq) + 2 E(s)  3 D(s) + 2 E3+(aq)
G+
OA
G +(aq) +
E(s)  no reaction
SOA
A3+(aq)
D2+(aq)
E3+(aq)
G+(aq)
+ 3 e- 
+ 2 e- 
+ 3 e- 
+ 1 e- 
A(s)
D(s)
E(s)
G(s)
SRA
Building Redox Tables
 So far we have been using examples where the oxidizing agents are metal ions
and the reducing agents are metal atoms. What else could gain or lose
electrons?
 Non-metal atoms I.e. Cl2(g) + 2e-  2 Cl-(aq) (Cl2(g) could act as a Reducing Agent)
 Non-metal ions I.e. 2 Br- (aq)  Br2(l) + 2 e- (2Br-(aq) could act as an Oxidizing Agent)
 Redox Table Trend


OA’s tend to be metal ions and non-metal atoms
RA’s tend to be metal atoms and non-metal ions
 Also, are there any entities that could act as both OA or RA?

Multivalent metals
Try Pg. 573 #14
Pg. 573 #14

Example 4: Use the following information to create a table of reduction ½ reactions
OA
Ag(s)
+
Br2(l)  AgBr(s)
Cl-
Br2(l)
OA
Ag(s)
+
Ag(s)
I2(s)  no evidence of reaction
I2(s)
OA
Cu2+(aq) +
I-(aq)  no redox reaction
Cu2+(aq)
OA
Br2(l) +
Cl-(aq)  no evidence of reaction
SOA
Cl2(g) +
2 e-
Br2(l) + 2 eAg+(aq) + 1 eI2(s) + 2 eCu2+(aq) + 2 e-
 2Cl-(aq)
 2Br-(aq)
 Ag(s)
 2I-(aq)
 Cu(s)
SRA
I-(aq)
Homework….
 Textbook:
 p.571 #1-3,6,7
 p.573 #12
 p.574 #15,17,20
 Diploma Questions:






Ex. MC P.25 #35, 37
Ex. MC P. 27 #39
Ex. MC P. 29 #42
MC 2013 #10, 12
NR 2013 #5
MC 2012 #12
Remember…
(Is Reduced)
(Is Oxidized)
Making a half reaction table
Try Pg. 573 #14
Pg. 573 #14

Example 4: Use the following information to create a table of reduction ½ reactions
OA
Ag(s)
+
Br2(l)  AgBr(s)
Cl-
Br2(l)
OA
Ag(s)
+
Ag(s)
I2(s)  no evidence of reaction
I2(s)
OA
Cu2+(aq) +
I-(aq)  no redox reaction
Cu2+(aq)
OA
Br2(l) +
Cl-(aq)  no evidence of reaction
SOA
Cl2(g) +
2 e-
Br2(l) + 2 eAg+(aq) + 1 eI2(s) + 2 eCu2+(aq) + 2 e-
 2Cl-(aq)
 2Br-(aq)
 Ag(s)
 2I-(aq)
 Cu(s)
SRA
I-(aq)
Predicting Redox Reactions (13.2b)
• Now that you know what redox reactions are, you will be responsible for
determining if a reaction will occur (is spontaneous) and if so, what the
reaction equation will be. How do we do this?
1. The first step is to determine all the entities that are present.
▫ Helpful reference: Table 6 pg. 575
▫ Remember: In solutions, molecules and ions behave
independently of each other.
▫ Example: When a solution of potassium permanganate is
slowly poured through acidified iron(II) sulfate solution.
▫ Does a redox reaction occur and what is the reaction equation?
Predicting Redox Reactions
2. The second step is to determine all possible OA’s and RA’s
▫
This is a crucial step!! Things to watch out for:

Combinations

(i.e. MnO4-(aq) is an oxidizing agent only in an acidic solution)

To indicate this draw an arc between the permanganate and
hydrogen ion

Species that can act as both OA and RA

Any lower charge multivalent metal i.e. Fe2+, Cu+, Sn2+, Cr2+

Water (H2O(l))

Label both possibilities in your list
Predicting Redox Reactions in Solution
• We can use a redox table to predict which reactions
will occur spontaneously.
• If we assume that collisions are completely random,
the strongest OA and the strongest RA will react.
• Tips for using the redox table:




Choose the strongest OA and RA present in your mixture
Read reduction half-reaction equations from left  right
Read oxidation half-reaction equations from right  left
Assume any substances not present in the table to be
spectator ions which do not need to be considered
Predicting Redox Reactions
3. The third step is to identify the SOA and SRA using the data booklet
SOA
SRA
3. The fourth step is to show the ½ reactions (from the redox table) and balance:
▫
SOA equation straight from table. SRA equation read from right to left
▫
▫
Are these equations balanced? Do the number of electrons lost = electrons gained
If not, multiply one or both equations by a number then add the balanced equations
Predicting Redox Reactions
5. The last step is to predict the spontaneity. Does the net
ionic equation represent a spontaneous or non-spontaneous
redox reaction??
If the SOA
above

Spontaneous
SRA??
If the
SRA
below
SOA
 Nonspontaneous
Predicting Redox Reactions #2
Could copper pipe be used to transport a hydrochloric acid solution?
1. List all entities
1. Identify all possible OA’s and RA’s
1. Identify the SOA and SRA
2. Show ½ reactions and balance
3. Predict spontaneity
Since the reaction is
nonspontaneous, it should be
possible to use a copper pipe to
carry hydrochloric acid
Disproportionation
• The redox reactions we have covered so far have one reactant (OA) which removes
electrons from a second reactant (RA) if a spontaneous reaction is to occur.
Although the OA and RA are usually different entities, this is not a requirement.
• A reaction is which a species is both oxidized and reduced is called
disproportionation (aka autoxidation or self oxidation-reduction)
▫ Occurs when a substance can act as either as oxidizing or reducing agent
▫ Example: Will a spontaneous reaction occur as a result of an electron transfer from
one iron(II) ion to another iron (II) ion?
▫ No! Using the redox table and spontaneity rule, we see that iron(II) as an oxidizing
agent is below iron(II) as a reducing agent, so the reaction is nonspontaneous
Disproportionation
• Example #2: Will a spontaneous reaction occur as a result of an electron
transfer from one copper(I) ion to another copper (I) ion? (see p. 828)
Cu+(aq) + 1 e-  Cu(s)
(Note: See pg. 578 Ex.2 for
another example.)
Cu+(aq)  Cu2+(aq) + 1 e2 Cu+(aq)  Cu2+(aq) + Cu(s)
▫ YES! Using the redox table and spontaneity rule, we see that copper(I) as an
oxidizing agent is above copper(I) as a reducing agent. Therefore, an aqueous
solution of copper(I) ions will spontaneously, but slowly, disproportionate into
copper(II) ions and copper metal.
Summary: Five Step method for
Predicting Redox Reactions
• Step 1: List all entities present and classify each
as a possible oxidizing agent, reducing agent, or
both. Do not label spectator ions.
• Step 2: Choose the strongest oxidizing agent as
indicated in a redox table, and write the equation
for its oxidation.
• Step 3: Choose the strongest reducing agent as
indicated in the table, and write the equation for
its oxidation
• Step 4: Balance the number of electrons
lost and gained in the half-reaction
equations by multiplying one or both
equations by a number. Then add the two
balanced half-reaction equations to obtain a
net ionic equation.
• Step 5: Using the spontaneity rule, predict
whether the net ionic equation represents a
spontaneous or nonspontaneous redox
reaction.
Homework…
• Textbook:
▫ p. 575 #25 a-c
▫ p. 579 #26
▫ p. 582 #4-7, 13
• Diploma Questions:
▫ MC 2011 #12, 16, 18, 23
▫ MC 2009 #15, 19, 25
Review:
• Data Booklet
▫ SOA top left
▫ SRA bottom right
▫ Reduction Half-Reactions (read
from left to right)
▫ Oxidization half-reactions (read
from right to left)
If the SOA
above

Spontaneous
SRA
If the
SRA
below
SOA
 Nonspontaneous
Review: Predicting Redox Reactions #2
Could copper pipe be used to transport a hydrochloric acid solution?
1. List all entities
1. Identify all possible OA’s and RA’s
1. Identify the SOA and SRA
2. Show ½ reactions and balance
3. Predict spontaneity
Since the reaction is
nonspontaneous, it should be
possible to use a copper pipe to
carry hydrochloric acid
Redox Reactions using Half-Reactions
13.2c
• Remember: a redox reaction includes BOTH an oxidation (RA) and a reduction
(OA).
• So far we have predicted redox reactions when the ½ reaction was provided to
us in the Redox table. But what if the table does not provide the half reaction?
• We can use our own knowledge to create the equation
Rules for Writing Half-Reactions
1.
2.
3.
4.
5.
Write an unbalanced ½ reaction showing formulas for reactants and products
Balance all atoms except H and O
Balance O by adding H2O(l)
Balance H by adding H+(aq)
Balance the charge by adding e- and cancel anything that is the same on both sides
For basic solutions only:
6. Add OH-(aq) to both sides to equal the number of H+(aq) present
7. Combine H+(aq) and OH-(aq) on the same side to form H2O(l). Cancel equal amounts of H2O(l)
from both sides.
Practicing Half-Reactions
• Copper metal can be oxidized in a solution to form copper(I) oxide.
What is the half-reaction for this process?
1. Balance all atoms except H and O
Cu(s)  Cu2O(s)
2Cu(s)  Cu2O(s)
1. Balance oxygen by adding water
2Cu(s) +H2O(l)  Cu2O(s)
1. Balance hydrogen by adding H+(aq)
2Cu(s) +H2O(l)  Cu2O(s) + 2H+(aq)
2. Balance charge by adding electrons
2Cu(s) +H2O(l)  Cu2O(s) + 2H+(aq) + 2 e-
Practicing Half-Reactions
• Chlorine is converted to perchlorate ions in an acidic solution. Write the half-reaction
equation. Is this half-reaction an oxidation or reduction? Cl2(g)  ClO4-(aq)
1.
Balance all atoms except H and O
1.
Balance oxygen by adding water
Cl2(g) + 8H2O(l)  2ClO4-(aq)
1.
Balance hydrogen by adding H+(aq)
Cl2(g) + 8H2O(l)  2ClO4-(aq) + 16H+(aq)
2. Balance charge by adding electrons
Cl2(g)  2ClO4-(aq)
Cl2(g) + 8H2O(l)  2ClO4-(aq) + 16H+(aq) + 14 e-
Cl2(g) + 8H2O(l)  2ClO4-(aq) + 16H+(aq) + 14 e-
OXIDATION
Predicting Redox Reactions by Constructing
Half-Reactions
SUMMARY
1. Use the information provided to start two half-reaction equations.
▫
Using the rules we just learned about half-reactions
2. Balance each half-reaction equation.
3. Multiply each half-reaction by simple whole numbers to balance electrons lost and gained.
4. Add the two half-reaction equations, cancelling the electrons and anything else that is
exactly the same on both sides of the equation.
5. (Remember to not if the solution is acidic or basic when writing half reactions.)
Predicting Redox Reactions by Constructing Half Reactions
•
•
Example: A person suspected of being intoxicated blows into this device and the alcohol in the
person’s breath reacts with an acidic dichromate ion solution to produce acetic acid
(ethanoic acid) and aqueous chromium(III) ions. Predict the balanced redox reaction
equation.
Create a skeleton equation from the information provided:
•
Separate the entities into the start of two half-reaction equations
•
Now use the steps you learned for writing half reactions
•
Now, balance the electrons lost and gained, and add the half reactions. Cancel the electrons and
anything else that is exactly the same on both sides of the equation.
Note: a similar
example is on p.581
(in basic solution)
Example #3 p.581
• Permananate ions and oxalate ions react in a basic
solution to produce carbon dioxide and
manganese(IV) oxide.
• MnO4-(aq) + C2O42-(aq)  CO2(g) + MnO2(s)
• Write a balanced redox equation for this reaction:
• 2[3 e- + 4 H+(aq) + MnO4-(aq)  MnO2(s) + 2H2O(l)]
• 3[ C2O42-(aq)  2 CO2(g) + 2e-]
• 8 H+(aq) + 2 MnO4-(aq) + 3 C2O42-(aq)  2 MnO2(s) +
4H2O(l) + 2 CO2(g)
Homework…
• Textbook:
▫ p. 581 # 31-32
▫ p. 582 #14-16
• Diploma (on a different paper):
▫ EX. #30, 31, 32
▫ MC 2011 #4
▫ MC 2001 #14, 20
▫ MC 2000 # 12