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Oxidation and Reduction
The chemical changes that occur when electrons are transferred between reactants are known as
_________________ - ________________ ______________. Oxidation reactions are the principal sources of
energy on earth. The combustion of gasoline in an automobile engine and the burning of wood in a fireplace
are oxidation reactions. So is the ‘burning’ of food by our bodies. All oxidation reactions are accompanied by
a reduction reaction. Oxidation-reduction reactions are also called _______________ reactions.
OXYGEN IN REDOX REACTIONS
Oxidation originally meant the combination of an element with ______________ to give oxides. For example,
when iron slowly turns to rust, it oxidizes to iron (III) oxide (Fe2O3). When carbon burns in air, it oxidizes to
carbon dioxide.
4Fe (s) + 3O2(g)  2 Fe2O3 (s)
C(s) + O2 (g)  CO2 (g)
Compounds can also be oxidized. Methane gas (CH4) burns in oxygen. It oxidizes to form oxides of carbon
and hydrogen.
CH4 (g) + 2O2 (g)  CO2 (g) + 2H2O (l)
Over the years, reduction has meant the ________ of oxygen from a compound. The reduction of iron ore to
metallic iron causes the removal of oxygen from iron (III) oxide. It is done by heating the ore with charcoal.
2 Fe2O3 (s) + 3C (s)  4Fe (s) + 3 CO2 (g)
As iron oxide loses oxygen, it is reduced to metallic iron. The term “reduction” refers to the fact that when a
metal oxide is reduced to the metal there is a considerable decrease in volume. In other words, the amount of
solid material has been reduced.
Oxidation and reduction occur ____________________. As iron (III) oxide is ______________ to iron
by losing oxygen, carbon is ________________ to carbon dioxide by gaining oxygen. No oxidation occurs
without reduction and no reduction occurs without oxidation. THIS IS TRUE OF ALL REDOX
REACTIONS.
ELECTRON TRANSFER IN REDOX REACTIONS
Today, chemists have extended the concepts of oxidation and reduction to include all transfers or shifts of
__________________. The advantage of the new definition is that it has much wider application than the old.
Hence, oxidation is _________________________________________________________________________
__________________________________________________________________________________________
Reduction is ______________________________________________________________________________
_________________________________________________________________________________________
A great way to remember this is the phrase: LEO the lion goes GER.
Some examples will show what this means. In the reactions between a metal and nonmetal, electrons are
transferred from atoms of the metal to atoms of the nonmetal. An ionic compound such as magnesium sulfide is
produced. (NOTE: We write the ionic compound as its ions with charge, rather than the compound itself.)
The net result of this reaction is a transfer of two electrons from a magnesium atom to a sulfur atom. The
magnesium atom loses two electrons and is ______________ to a magnesium ion. Simultaneously the sulfur
atom gains two electrons and is _________________ to a sulfide ion.
Oxidation: Mg  Mg +2 + 2e- (loss of electrons)
Reduction: S + 2e-  S -2 (gain of electrons)
The substance in a redox reaction that donates electrons is a ______________ ________. By losing electrons,
magnesium ____________ sulfur. Magnesium is a reducing agent.
The substance in a redox reaction that accepts electrons is an _____________ ________. By accepting
electrons, sulfur _____________ magnesium. Sulfur is an oxidizing agent.
Example: What is oxidized and what is reduced in this single replacement reaction?
2AgNO3 (aq) + Cu (s)  Cu(NO3)2 (aq) + 2Ag (s)
To begin solving this question, rewrite the equation by showing the ionic compound as their ions.
Now let’s see. What happened in this reaction? Two electrons have been transferred from a copper atom to
two silver ions.
Oxidation:
Reduction:
The Ag gains electrons and is ___________________. The Cu loses electrons and is ____________________.
Hence Ag is the __________________ agent, and Cu is the __________________ agent.
Practice:
1. Identify each of the following changes as either oxidation or reduction.
a) I2 + 2 e-  2Ib) K  K+ + ec) Fe2+  Fe3+ + ed) Ag+ + e-  Ag
2. Identify what is oxidized and what is reduced in the following processes.
a) 2Br- + Cl2  Br2 + 2Clb) 2Ce + 3Cu2+  3Cu + 2Ce3+
c) 2Zn + O2  2ZnO
3. Identify the oxidizing agent and the reducing agent in each of the following reactions.
a) Mg + I2  MgI2
b) 2Na + 2H+  2Na+ + H2
c) H2S + Cl2  S + 2HCl
4. Determine what is oxidized and what is reduced in each reaction. Identify the oxidizing agent and
reducing agent in each case.
a) 2Na(s) + S(s)  Na2S (s)
b) 2K (s) + Cl2 (g)  2KCl (s)
c) 4Al (s) + 3O2 (g)  2Al2O3 (s)
ASSIGNING OXIDATION NUMBERS
Oxidation numbers are a book keeping concept devised by chemists. An oxidation number is a ___________
or ___________ number assigned to an atom according to a set of arbitrary rules. Complete redox equations
can be balanced by using oxidation number charges. As a general rule, an oxidation number is the charge of
a monatomic ion.
For example  NaCl
This compound of sodium chloride consists of Na1+ ions and Cl1- ions. Thus, the oxidation number of sodium is
+1 and of chlorine is -1. Notice that when writing oxidation numbers, the sign is put before the number.
Sodium has an ionic charge of 1+ and an oxidation number of +1. In calcium fluoride, CaF2, the oxidation
number of calcium is ___ and that of fluorine, ___.
Because water is a molecular compound, no ionic charges are associated with its atoms. However, oxygen is
considered to be reduced in the formation of water. In water, the shared electrons in the bond are shifted closer
to oxygen and away from hydrogen. Imagine that the two electrons contributed to the hydrogen atoms were
completely transferred to the oxygen. The charges that would result from this transfer are the oxidation
numbers of the elements. The oxidation number of oxygen is -2 and +1 for each hydrogen.
Rules for determining oxidation numbers:
1. The oxidation number of a monatomic ion is equal to the charge on the ion.
Ie. The oxidation number of Br1- is ____, iron (III) has an oxidation number of __
Ex) Na3N(s) --> 3Na+(aq) + N3-(aq)
2. The oxidation number for metals in an ionic compound is just their ionic charge.
Ie. The oxidation number of Ca in CaBr2 is _____, the oxidation number for Fe in
Fe2O3 is _______.
Ex) 2NaBr(aq) + Hg(NO3)2(aq) --> 2NaNO3(aq) + HgBr2(s)
3. The oxidation number of hydrogen in most of its compound is +1. The exception is when hydrogen
bonds to metals to make a metal hydride.
Ie. In NaH, H would be have an oxidation number of ____. This is because hydrogen attracts electrons
more strongly than does the metal atom.
Ex) HBr(aq) + NaOH(aq) --> NaBr(aq) + H2O(l)
4. The oxidation number of oxygen in a compound is always -2, except in peroxides, such as hydrogen
peroxide (H2O2) where it is ____. When it is bonded to fluorine, the only element more electronegative than
oxygen, the oxidation number of oxygen is +2.
Ex) AlPO4 + NaNO3 --> Al(NO3)3 + Na3PO4
5. The oxidation number of an uncombined atom is zero. This is true for elements that exist as diatomic
molecules such as O2, Cl2, H2, etc.
Ie. The oxidation in potassium metal, K, and the nitrogen atoms in nitrogen gas, N2, is zero.
Ex) Mg(s) + O2(g) --> MgO(s)
6. For any neutral compound, the sum of oxidation numbers of the atoms in the compound must equal
zero. Notice how the oxidation numbers add up to zero in the following examples.
Ie.
NaCl
CaBr2
H2S
7. For a polyatomic ion, the sum of the oxidation numbers must equal the ionic charge of the ion.
Ie.
NH4+
SO32-
CO32-
Practice: Determine the oxidation number of the boldface element in these formulas.
a) P2O5
f) AlPO4
b) NaClO4
g) HNO2
c) Na2Cr2O7
h) AsO43-
d) H3PO4
i) CrO42-
e) S2O3 2-
Oxidation Number Changes:
An increase in the oxidation number of an atom signifies oxidation.
A decrease in the oxidation number of an atom signifies reduction.
2AgNO3 (aq) + Cu (s)  Cu(NO3)2 (aq) + 2Ag (s)
In this reaction, the oxidation number of silver decreases from +1 to 0. This is _____________: silver is
______________ from Ag +1 to Ag 0.
On the other hand, copper is ______________ in this reaction from Cu 0 to Cu +2. These results agree with
those obtained by analyzing the reaction by using electron transfer.
Example: Use the change in oxidation number to identify which elements are oxidized and reduced in each of
these reactions.
a) Cl2 (g) + 2HBr (aq)  2HCl (aq) + Br2 (l)
b) C (s) + O2 (g)  CO2 (g)
Identifying Redox Equations:
In general, all chemical reactions can be assigned to one of two classes. In oxidation-reduction reactions
electrons are transferred from one reacting species to another. In all other reactions electrons are not
transferred. A majority of the reactions presented prior to this unit do not involve electron transfer and are not
redox equations. Double replacement reactions and acid base reactions are not redox reactions. Many single
replacement reactions, combination reactions, decomposition reactions and combustions reactions are redox
reactions.
So, how can you determine if a reaction is a redox reaction? Well, simply use oxidation numbers. If the
oxidation number of an element in a reacting species changes then that element has undergone either oxidation
or reduction.
Practice:
1. Cl2 (g) + 2NaBr (aq)  2NaCl (aq) + Br2 (g)
2. PbCl2 (s) + K2SO4 (aq)  2KCl (aq) + PbSO4 (s)
3. 2K (s) + 2H2O (l)  2KOH (aq) + H2 (g)
4. BaCl2 (aq) + 2KIO3(aq)  Ba(IO3)2 (s) + 2KCl(aq)
5. H2CO3  H2O + CO2
6. Mg(s) + Br2 (l)  MgBr2 (s)
7. 2Sb (s) + HNO3 (aq)  Sb2O5 (s) + NO (g) + H2O (l)
Balancing Redox Equations:
Simple Redox equations can be balanced by inspection or by trial and error. Many other oxidation-reduction
reactions are too complex, however, to balance in this way. Fortunately, two systematic methods are available.
These are the oxidation-number change method and the half reaction method. These methods are based on
the fact that the total number of electrons gained in reduction must equal the total number of electron
lost in oxidation.
With the oxidation number change method, a redox equation is balanced by comparing the increases and
decreases in oxidation numbers. To use this method, start with a skeleton equation for a redox reaction. We
will use the reduction of iron ore as an example.
Fe2O3 (s) + CO (g)  Fe (s) + CO2 (g)
Step 1: Assign oxidation numbers to all the atoms in the equation. Write the number above the appropriate
atoms.
Fe2O3 (s) + CO (g)  Fe (s) + CO2 (g)
Note that the oxidation number is stated as the charge per atom. In Fe2O3 the total positive charge is +6 but the
oxidation number of Fe is +3.
Step 2:



Identify which atoms are oxidized and which are reduced. In this reaction:
Fe decreases in oxidation number from +3 to 0, a change of -3. Therefore iron is ________________.
C increases in oxidation number from +2 to +4, a change of +2. Thus carbon is ________________.
Oxygen does not change in oxidation number.
Step 3: Use a line to connect the atoms that undergo oxidation and those that undergo reduction. Write the
oxidation number change at the midpoint of each line.
+3 -2
+2-2
0
+4 -2
Fe2O3 (s) + CO (g)  Fe (s) + CO2 (g)
Step 4: Make the total increase in oxidation number equal to the total decrease in oxidation number by using
appropriate coefficients. In this example,
 the oxidation number increase should be multiplied by ___ and
 the oxidation number decrease should be multiplied by ___.
 This gives an increase of +6 and a decrease of -6.
 Now use coefficients in the formula to express the above multiplications.
+2
+3 -2
+2 -2
0
+4 -2
Fe2O3 (s) + CO (g)  Fe (s) + CO2 (g)
-3
Note: Fe2O3 does not need a coefficient because
the formula already indicates that there are 2 Fe.
Step 5: Finally, check to be sure that the equation is balanced for both atoms and charge. If necessary, the
remainder of the equation is balanced by inspection.
Fe2O3 (s) + 3CO (g)  2Fe (s) + 3CO2 (g)
Now you try one
Balance this redox equation by using the oxidation number change method. (Remember to complete other
balancing by inspection!)
K2Cr2O7 (aq) + H2O (l) + S(s)  KOH (aq) + Cr2O3 (aq) + SO2 (g)
Some can be balanced using fractions:
Fe + Br2  FeBr3
Fe + Br2  FeBr3
Practice:
Balance the following redox equations: (you may have to use fractional coefficients for some!)
a)
HNO3 (aq) +
b)
FeCl3 +
c)
Al +
H2S (aq) 
Zn 
Cl2 
ZnCl2 +
S (s) +
NO (g) +
H2O (l)
Fe
AlCl3
You can use
fractions as
coefficients
to balance
d)
KClO3 
e)
HCl +
f)
SnCl4 +
KCl +
HNO3 
Fe 
O2
HOCl +
SnCl2 +
NO +
FeCl3
H2O