Transcript Redox Reaction

-Types of Redox Reactions
- Balancing Redox Reactions
Mr. Shields
Regents Chemistry
U14 L02
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Types of Redox Reactions
We’ve talked about 5 different types of chemical reactions
Before:
1) Synthesis
2) Decomposition
3) Combustion
4) Single Replacement
5) Double Repalcement
The first 4 of these reaction types are examples of Redox
Reactions. Why isn’t double replacement an example of a
Redox Reaction?
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In double replacement two ions simply replace each other
In the compounds undergoing reaction. There is no REDOX
Occuring.
For example:
BaCl2 + AgNO3  AgCl + Ba(NO3)2
What’s the oxidation nos. of the reactants and Products
In the equation above?
Ba+2
Cl-1
Ag+1
N+5
O-2





Ba+2
Cl-1
Ag+1
N+5
O-2
Has any oxidation or reduction taken place?
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Synthesis
In Synthesis, elements or compounds combine to form
One product.
Example: N2 + 3H2  2NH3
This is a Redox Reaction. What is being Oxidized and
What’s being Reduced?
N0 and H0 change to N-3 and H+1. Nitrogen has been reduced
And Hydrogen has been oxidized
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Decomposition
In Decomposition a compound is broken down into simpler
Compounds and/or elements.
Example: 2HgO  2Hg + O2
This too is a Redox reaction. What is being Oxidized &
What’s Reduced?
Hg+2 and O-2 become Hg0 and O0. Hg+2 has been reduced
And O-2 has been oxidized.
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Combustion
In Combustion a compound reacts with oxygen to form
several compounds, if combustion is complete the products
Are always water and carbon dioxide.
Example: CH4 + 2O2  2H20 + CO2 + Heat
Again, this is a Redox reaction. What’s being Oxidized &
What is being Reduced?
C-4 and O0 become C+4 and O-2. Carbon-4 has been oxidized
And O0 has been reduced.
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Single Replacement
In Single Repalcement rxns a free element replaces another
Element that is combined in a compound.
Example: Zn + Cu(NO3)2  Zn(NO3)2 + Cu
Since this is a Redox rxn what is being Oxidized & Reduced?
Zn0 and Cu+2 become Zn+2 and Cu0 . Zinc (0) has been
Oxidized and Cu (+2) has been reduced.
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Remember when we’re discussing single replacement rxns
We need to consider what their position is relative to one
Another on TABLE J
Metals higher on the
List will replace metals
Lower on the list.
Most easily oxidized
Metals higher on the
List are more reactive
And lastly, metals
Higher on the list are
More easily oxidized.
Most easily reduced
Less easily oxidized
Leasteasily reduced
Notice that halogens accept electrons so higher on the list
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means more easily reduced!
Problem: in the following reaction
Sr + FeCL2  SrCl2 + Fe
What gets oxidized?
Sro
What gets reduced?
Iron +2
What is the oxidizing agent?
Iron +2
What is the reducing agent?
Sro
Write the Oxidation & Reduction half cell rxns
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Balancing Redox Eqs
We’re now going to return to a topic we’ve discussed
Before… BALANCING CEHEMICAL EQUATIONS
For example: Fe + O2  Fe2O3
1) 1st we count atoms on the reactant side and then on the
Product side
#R
1
2
atom #P
Fe
2
O
3
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Balancing Redox Eqs
2) We then pick one atom to balance, let’s say oxygen
3x
#R
1
2
atom #P
Fe
2
O
3
Then see what we’ve changed:
x2
Fe + 3O2  2Fe2O3
3) After we’ve adjusted our table we then balance the next
atom, Fe
4x
#R
1
6
atom #P
Fe
4
O
6
balanced
4Fe + 3O  2Fe O
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Balancing Redox Rxns
Many reactions can be balanced “visually” or
using the RAP Method we just reviewed.
However, many redox reactions are difficult to balance
Using these methods.
Consider the difficulty of balancing this REDOX equation:
S + HNO3  SO2 + NO + H2O
We could balance it with RAP tables but it would be difficult
To more easily balance this equation we need to apply
A method that will utilize oxidation numbers.
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Balancing Redox Rxns
There are some rules you need to follow to balance Redox
reactions. By balancing reactions this way we will not only
Balance the equation but we will BALANCE THE HALF CELLS
1) The number of e- lost in an oxidation process must equal
the number of electrons gained in the reduction process
2) Assign oxidation numbers to all atoms in the chem. Eq.
3) Identify which reactants are oxidized & which are reduced
0
+1 +5 -2
+4 -2
+2 -2
+1 -2
S + HNO3  SO2 + NO + H2O
oxidized
reduced
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Balancing Redox Rxns
4) Visually balance any atoms that change from ions to atoms
or vice versa
EX:
Cr + Cl2  CrCl3
becomes
Cr + 3Cl2  2CrCl3
And
NH3 + O2  N2 +H20
becomes
2NH3 + O2  N2 + 2H2O
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Balancing Redox Rxns
5) Connect the atoms that change oxidation number by a line
6) Write the change in electrons on the line
Ox. -4 (e-)
0
+1 +5 -2
S + HNO3
+4 -2 +2 -2 +1 -2
 SO2 + NO + H2O
Red +3 (e-)
Note that these brackets define the 2 half cell reactions
S0  S+4 + 4e-
and
N+5 + 3e-  N+2
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Balancing Redox Rxns
7) Multiply the change in electrons by a number that makes
the # of ox. electrons equal to the # of red. electrons
Ox. -4 (e-) x 3
0
+1 +5 -2
+4 -2
+2 -2
+1 -2
S + HNO3  SO2 + NO + H2O
Red +3 (e-) x4
8) These nos. become the coefficients in the eq.
3S + 4HNO3  3SO2 + 4NO + H2O
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Balancing Redox Rxns
8) Balance the rest of the equation by visual inspection
or RAP tables
3S + 4HNO3  3SO2 + 4NO + 2 H2O
Need 4 H on product side
Increase H to 4
Is oxygen also balanced?
Yes. There are 12 O on each side of the Eq.
This equation is now balanced.
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Problem: Balance the following equation!
Ag + HNO3  AgNO3 + NO + H20
Solution:
1)
Ox -1 e- x3
0
+1 +5 -2
+1 +5 -2
+2 -2
+1 -2
Ag + HNO3  AgNO3 + NO + H20
Red. +3 e- x1
2)
3Ag + HNO3  3AgNO3 + NO + H20
Visually or using RAP tables Balance the Rest of the Eq.
3)
3Ag + 4HNO3  3AgNO3 + NO + 2H20
1) I Need 4 N
Is “O” OK?
2) SO… I need 4 H
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Notice that the half cell reactions are now also
=
balanced
Ox -1 e- x3
0
+1 +5 -2
+1 +5 -2
+2 -2
+1 -2
Ag + HNO3  AgNO3 + NO + H20
Red. +3 e- x1
3Ag + 4HNO3  3AgNO3 + NO + 2H20
Unbalanced half cell reactions:
Ag0  Ag+1 + 1eN+5 + 3e-  N+2
Balanced half cell reactions:
3Ag0  3Ag+1 + 3e1 N+5 + 3e- 1 N+2
# of e- lost
And gained
Are equal
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Conservation of Mass & Charge
Remember !!
In ALL chemical reactions there is not only
conservation of mass
- Equal number of atoms on both sides
There must also be conservation of charge
- Number of e- lost = Number of e- gained
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